Environ. Sci. Technol. 2010, 44, 8529–8534
Reaction Mechanism for Chlorination of Urea ERNEST R. BLATCHLEY, III, * AND MINGMING CHENG School of Civil Engineering and Division of Environmental & Ecological Engineering, Purdue University, West Lafayette, Indiana 47907-2051, United States
Received July 14, 2010. Revised manuscript received September 27, 2010. Accepted October 3, 2010.
Experiments were conducted to elucidate the mechanism of the reaction between free chlorine and urea. In combination with findings of previous investigations, the results of these experiments indicate a process by which urea undergoes multiple N-chlorination steps. The first of these steps results in the formation of N-chlorourea; this step appears to require Cl2 to proceed and is the overall rate-limiting step in the reaction for conditions that correspond to most swimming pools. N-Chloroureathenappearstoundergofurtherchlorinesubstitution; the fully N-chlorinated urea molecule is hypothesized to undergo hydrolysis and additional chlorination to yield NCl3 as an intermediate. NCl3 is hydrolyzed to yield NH2Cl and NHCl2, with subsequent decay to stable end products, including N2 and NO3-. Conversion of urea-N to nitrate is pHdependent. The pattern of nitrate yield is believed to be attributable to the fact that when urea serves as the source of reduced-N, entry into the reactions that describe chlorination of ammoniacal nitrogen is through NCl3, whereas when NH3 is the source of reduced-N, entry to these reactions is through NH2Cl.
Introduction Swimming is second only to walking as a form of exercise in the United States, with broad participation among all age and income groups (1). Aquatic activities are widely recognized as being beneficial in terms of cardiovascular health and fitness and are used as therapy to treat a wide range of medical conditions. In the majority of cases, swimming activities take place in public or private pools where chlorine is used as a disinfectant and oxidant. Important limitations associated with chlorination of pools have been identified, including formation of disinfection byproducts (DBPs) (2, 3). DBP formation in swimming pools has emerged as an important area of study (4). Among indoor swimming pool facilities, the formation of volatile DBPs is of particular concern because of the association of some of these compounds with human health problems, including promotion of asthma (5), increased incidence of rhinitis and hay fever (6), and skin (contact dermatitis) and eye irritation (7). Relative to the general population, people who have experienced high levels of swimming pool DBP exposure (e.g., lifeguards and elite swimmers) tend to be most likely to display these symptoms. LaKind et al. (8) present an excellent overview of these issues. Trichoramine (NCl3) has been the focal point of most in* Corresponding author phone: 765-494-0316; e-mail: blatch@ purdue.edu. 10.1021/es102423u
2010 American Chemical Society
Published on Web 10/21/2010
vestigations related to human health effects of exposure to products of chlorination in swimming pools. Many of the volatile DBPs that have been identified in pools contain nitrogen (3). Known precursors to the formation of volatile N-DBPs in pools include urea, uric acid, creatinine, and several amino acids. By mass, urea is the predominant source of organic-N in human sweat and urine (9), which are the most likely sources of DBP precursors in pools. The chemistry of reactions between chlorine and urea was first examined by Chattaway (10), who suggested that chlorine reacts with urea to yield N,N′-dichlorourea as an intermediate; N2 and NCl3 were identified as products of this reaction. On the basis of experiments at pH ≈ 7.2, Palin (11) identified NCl3 formation from urea-chlorine reactions as being attributable to urea hydrolysis. Samples (12) argued that the near-neutral pH conditions and low microbial concentrations that exist in swimming pools would rule out urea hydrolysis, and that reactions between chlorine and urea in pools would involve some other mechanism. On the basis of spectrophotometric measurements, titrations, and conjecture, Samples hypothesized a reaction mechanism in which urea underwent sequential chlorine substitution to yield N,N,N′,N′-tetrachlorourea, which was then hydrolyzed in the presence of chlorine to yield NCl3. The reaction was promoted under low pH conditions, and the first chlorine substitution was reported to be slower than subsequent steps. The results of previous investigations indicate that NCl3 is a product of urea chlorination. Conditions of low pH appear to promote this process, but some uncertainty remains regarding the mechanism of this reaction, including an explanation for the pH-dependence and the fate of nitrogen in these reactions. The objective of the experiments described herein was to further elucidate the mechanism of reactions between free chlorine and urea. Experiments were conducted to quantify the dynamics of reactions that resulted from addition of free chlorine and urea under different conditions of solution chemistry and temperature, and to further elucidate the kinetics of the initial step in the reaction process (Nchlorination), which appears to limit overall reaction progress.
Materials and Methods Reactions were performed in headspace-free, gas-tight, glass containers that had been soaked in a free chlorine solution to remove chlorine-demanding substances and then rinsed with deionized water. Reaction mixtures were introduced to reaction vessels and then submerged in a water bath for temperature control. Reaction vessels were maintained in the dark during experiments to eliminate contributions from photochemical processes. Samples were removed periodically for analysis. Vessels were returned to the water bath after sample collection. The headspace created by sample removal was sufficiently small that equilibration of volatile compounds with the headspace allowed only trace quantities of volatile DBPs to be lost from solution. Reacting solutions were prepared by addition of carbonate buffer (10-3 M) to deionized water; pH adjustment was accomplished by addition of a strong acid (H2SO4) or strong base (NaOH). Urea (Sigma-Aldrich) was added gravimetrically to produce a stock solution, which was then diluted to achieve target concentrations for each experiment. Free chlorine was added in the form of an aqueous stock solution of NaOCl (Sigma-Aldrich). The free chlorine stock solution was standardized periodically using DPD/FAS titration. A matrix of reaction conditions was examined; matrix variables included temperature (T) ) 20, 25, and 30 °C, and VOL. 44, NO. 22, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. Free chlorine (top) and NCl3 (bottom) concentration as a function of time and pH for T ) 25 °C and Cl:P ) 1.0. Initial urea concentration ) 5.0 × 10-5 M. Spline fits are included to illustrate trends in the data.
FIGURE 2. Free chlorine (top) and NCl3 (bottom) concentration as a function of time and temperature for pH ) 7.5 and Cl:P ) 1.0. Initial urea concentration ) 5.0 × 10-5 M. Spline fits are included to illustrate trends in the data.
pH ) 6.5, 7.5, and 8.5. For each combination of T and pH, solutions were prepared at free chlorine:precursor (Cl:P) molar ratios of 0.5, 1.0, 1.5, 2.0, 2.5, and 3.0. Urea was added at an initial concentration of 5.0 × 10-5 M (100 µM N). Samples were collected for analysis after 0-24 h of chlorination. Urea concentration is highly variable in pools but generally is 1-2 orders of magnitude lower than in the experiments described herein (Table SI-1, Supporting Information). Free chlorine concentration in pools generally ranges from 1 × 10-5 M to 7 × 10-5 M. Free chlorine was analyzed using a DPD spectrophotometric method. Inorganic combined chlorine was measured using membrane introduction mass spectrometry (MIMS) (13). Urea was quantified using the method of Prescott and Jones (14). Nitrate and nitrite were quantified by ion chromatography. A stopped-flow device was used to characterize the kinetics of the first step in the reaction, under room temperature conditions (23 °C). Absorbance was used as a means of quantifying reaction dynamics in real time. Absorbance spectra for reactants and pure compounds that were known or suspected to form as intermediates or products were measured using a Perkin-Elmer Lambda 20 UV-visible scanning spectrophotometer. This same spectrophotometer was used in time-course measurements of absorbance at a fixed wavelength for some reacting solutions.
Cl:P molar ratio of 1.0 at various pH values and T ) 25 °C; similar trends were evident at other Cl:P ratios. Figure 1 indicates that the dynamics of free chlorine and NCl3 are sensitive to pH in the range 6.5-8.5. Urea is a weak base (pKa ≈ 0.1); therefore, it is unlikely that the pHdependence illustrated in Figure 1 could be attributable to the acid-base behavior of urea; however, this pH range, which is representative of conditions in most pools, can have a profound effect on chlorine speciation. Therefore, it is hypothesized that the pH-dependent changes in chlorine consumption and NCl3 formation illustrated in Figure 1 are attributable to shifts in free chlorine speciation. A decrease of pH would result in a shift of free chlorine toward HOCl and Cl2, but the data in Figure 1 are insufficient to differentiate between the contributions of these two forms of free chlorine. Note that the reaction between free chlorine and urea is relatively slow at near-neutral pH and temperatures that are representative of recreational water applications. For example, at pH ) 7.5, the half-life for free chlorine was greater than one day. Figure 2 illustrates the dynamic behavior of free chlorine and NCl3 as functions of temperature for pH ) 7.5 and Cl:P ) 1.0. The rate of chlorine consumption in the presence of urea increased with temperature, as did the rate of NCl3 formation. At pH ) 7.5, the rate of chlorine consumption by urea more than doubled when the temperature increased from 20 to 30 °C. NCl3 is an intermediate in this process. The rates of NCl3 formation and decay both increase with temperature. The range of temperatures illustrated in Figure 2 is representative of the range of operating conditions normally encountered in pools. Clearly, the rates of chlorine consumption and NCl3 formation from urea will be influenced by pool pH and temperature.
Results and Discussion Effects of pH and Temperature. The dynamics of reactions between free chlorine and urea were first characterized by time-course measurements of residual chlorine concentrations. Figure 1 illustrates the dynamic behavior of free chlorine and NCl3 when free chlorine and urea were introduced at a 8530
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TABLE 1. Concentrations (µM) of N-Compounds in Solution after 24 h of Chlorination at T = 25°C, Cl:P = 3.0, and with an Initial Urea Concentration of 5.0 × 10-5 M (100 µM N)
FIGURE 3. Residual chlorine concentration as a function of time and temperature for pH ) 6.5, T ) 30 °C, and Cl:P ) 2.0. Inorganic chloramine concentrations were measured by MIMS. Initial urea concentration ) 5.0 × 10-5 M. Spline fits are included to illustrate trends in the data. Note: inorganic chloramine concentrations refer to the left axis; free chlorine concentrations refer to the right axis. NCl3 has long been recognized as a product of urea chlorination (3, 10-12). The dynamic behavior of NCl3 is influenced by reactions that lead to its formation, as well as reactions that result in its decay by oxidation. In swimming pools, NCl3 is generally present with free chlorine. Obviously, free chlorine reacts with urea to produce NCl3. However, free chlorine will also oxidize NCl3; the kinetics and mechanism of this reaction have been described in the literature (15). Therefore, NCl3 exists as an intermediate in the reactions between chlorine and urea. As described above, the conditions that promote NCl3 formation (low pH and high temperature) are the same as those that promote loss of free chlorine. These conditions also appear to promote reactions that lead to loss of NCl3, presumably through subsequent NCl3 oxidation. NCl3 formation dictates that the other inorganic chloramines will also be present (15). Figure 3 illustrates the timedependent behavior of inorganic chloramine concentration and free chlorine for chlorination of urea at pH ) 6.5, T ) 30 °C, and Cl:P molar ratio of 2.0 (see also Figure SI-1, Supporting Information). The trends in NCl3 and free chlorine behavior in Figure 3 (and Figure SI-1) are similar to those illustrated in Figures 1 and 2, where experiments were performed at a lower Cl:P ratio. However, the pattern of inorganic chloramine behavior illustrates an interesting trend. In particular, NCl3 was present at higher concentrations than either NH2Cl or NHCl2. This pattern is fundamentally different than the distribution of inorganic chloramines that results from chlorination of ammonia. In the chlorination of ammonia-N, chlorine substitution and hydrolysis reactions play important roles. Chlorine substitution reactions result in replacement of aminohydrogen by +1-valent chlorine (Cl+); these reactions are common in compounds that contain amine groups, and have been widely reported as being the first reactions between free chlorine and many amines. Hydrolysis reactions can lead to cleavage of an amine group from the parent molecule. It is likely that formation of NCl3 by chlorination of urea will involve both of these reactions. One explanation for the data presented in Figure 3 and Figure SI-1 is that the amine group is being cleaved from the parent molecule only after it undergoes complete chlorine substitution. This would yield a dichlorinated amino radical, which presumably would react with free chlorine to produce NCl3. The formation of any of the inorganic chloramines opens up the entire mechanism that describes reactions
pH
urea
NH2Cl
NHCl2
NCl3
NO3-
N recovery (%)
6.5 7.5 8.5
65.1 70.1 63.3
0.381 0.266 0.188
1.36 0.964 0.796
2.96 1.85 1.22
6.23 5.01 4.44
76.1 78.2 70.4
between free chlorine and ammonia (15). However, rather than entering this mechanism through NH2Cl formation, it appears that this reaction sequence leads to the process as a result of NCl3 production. The formation of NH2Cl and NHCl2 would then be expected to follow as a result of NCl3 hydrolysis. Table 1 provides a summary of the fate of nitrogen in these experiments. Because the method of Prescott and Jones (14) does not differentiate between urea and its N-chlorinated forms, the values reported for urea should be interpreted as the sum of the concentrations of urea and its N-chlorinated forms. All data presented in Table 1 correspond to the distribution of N-compounds after 24 h of reaction time. Collectively, several conclusions can be drawn from the data presented in Table 1: • The reactions between free chlorine and urea are slow to yield inorganic chloramines and nitrate. The vast majority of N in the system remained as urea after 24 h of reaction time. As described previously, the urea concentration used in these experiments was roughly 1 order of magnitude greater than the urea concentrations that have been measured in pools. Therefore, urea-chlorine reactions in pools would be expected to proceed at slower rates than were observed in these experiments. As such, urea represents a reservoir for formation of NCl3 and other compounds in pools. • The compounds listed in Table 1 account for the majority of N in the system after 24 h of reaction time. Given that the inorganic chloramines are formed in the reactions between free chlorine and urea, it is likely that a substantial fraction of the urea-N is converted to molecular nitrogen (N2), as N2 is a known, stable product of oxidation of ammoniacal nitrogen (15). N2 formation was not quantified in this research. • Urea is an efficient precursor to NCl3 formation, although the reactions that lead to NCl3 formation are relatively slow. • Chlorination of urea leads to NO3- formation in pools. Nitrate-N accounted for 12-18% of the urea-N that was lost due to these reactions, depending on pH. NO3- is a stable, nonvolatile compound and is expected to accumulate in pools where the presence of free chlorine will prevent reduction to NO2-. Chlorination of Urea at Cl:P < 1. Experiments were conducted at Cl:P < 1 (i.e., excess urea, relative to free chlorine) as a means of isolating the first step in the reaction process, which was believed to be N-chlorination of the urea molecule. Absorbance spectra for urea and N-chlorourea were measured as a means of determining the ability to differentiate these compounds spectroscopically. The urea absorbance spectrum was measured using aqueous solutions of urea at concentrations ranging from nominally 0.01-10 M (Figure SI-2). For reference purposes, the value of the molar absorptivity of urea at λ ) 245 nm was 0.0172 M-1 · cm-1. An aqueous solution of N-chlorourea was prepared by addition of free chlorine and urea to aqueous solution at Cl:P , 1, so as to limit N-chlorination to a single step. Methods were not available to confirm the structure of the product of this reaction. However, the reaction leading to formation of this compound was conducted under low Cl:P molar ratio. The free chlorine originally present in solution was lost and was consistent with the corresponding increase in combined VOL. 44, NO. 22, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 4. The N-chlorourea yield as a function of pH after 4 h of reaction with free chlorine. Initial conditions were [urea] ) 3.00 mM, free chlorine concentration ) 1.76 × 10-4 M, thereby yielding an initial Cl:P molar ratio of 1:17. The horizontal red line indicates the initial molar concentration of free chlorine (right axis). chlorine measured by DPD/KI spectrophotometric analysis. Figure SI-3 illustrates the absorbance spectrum for Nchlorourea. The data presented in Figures SI-2 and SI-3 indicate that N-chlorourea and urea can be differentiated spectroscopically, and that the N-chlorourea concentration can be monitored by measurements of absorbance at wavelengths in the range of 240-250 nm. For reference purposes, the value of molar absorptivity for N-chlorourea at λ ) 245 nm was measured as 245 M-1 · cm-1, which is similar to the molar absorptivity reported by Samples (12) of 240 M-1 · cm-1 (λ ) 245 nm). Figure 4 illustrates the N-chlorourea yield as a function of pH after 4 h of reaction time. For all data illustrated in Figure 4, Cl:P ) 1:17 was imposed on the solution at t ) 0. The N-chlorourea concentration was then measured after 4 h of reaction time, based on measurements of A245. The data in Figure 4 indicate that formation of N-chlorourea is favored by conditions of low pH. At pH ) 3.0, N-chlorination was brought to completion in less than 4 h. Under these conditions, virtually all free chlorine initially present in solution was bound to the urea molecule. At low pH, a shift in free chorine distribution will promote formation of HOCl and Cl2. However, the data presented in Figure 4 are more consistent with Cl2 as the reactive form of free chorine than HOCl, as the N-chlorourea yield continued to increase as pH decreased well below the pKa of HOCl. These same initial conditions were replicated to allow time-course evaluation of the concentrations of N-chlorourea and free chlorine. Figure 5 illustrates these dynamics for conditions of pH ) 7.5 and pH ) 3.0. These conditions were selected to allow for examination of the effects of pH on N-chlorination of urea. For the pH ) 3.0 condition, Nchlorination proceeded (essentially) to completion rapidly. After completion of the reaction, where N-chlorourea would have been present in solution in the presence of excess urea and virtually no free chlorine, the solution chemistry appeared to remain essentially unchanged for the duration of the experiment. Therefore, N-chlorourea appears to be stable on a time-scale of at least the duration of this experiment (ca. 10 h) in the absence of free chlorine. By contrast, in the pH ) 7.5 experiment, N-chlorourea appeared to form as an intermediate. The pattern of the A240 signal at pH ) 7.5 was inconsistent with formation of N-chlorourea alone. Decay of the N-chlorourea signal (top panel) coincided with loss of free chlorine. This suggests that N-chlorourea will react with free chlorine to yield one or more compounds that do not absorb strongly at 240 nm. The ratio of measured changes in free chlorine concentration (∆Cl) to the change in A240 8532
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FIGURE 5. Dynamic behavior of N-chlorourea (top panel, as measured by A240) and free chlorine (bottom panel, as measured by DPD/FAS titration) at pH ) 7.5 and pH ) 3.0. Initial conditions were otherwise identical to those described in Figure 4. signal (∆A240) were different at these two pH values. For example, at t ) 30 min and pH ) 3.0, ∆Cl:∆A240 ) 4690, whereas at pH ) 7.5, ∆Cl:∆A240 ) 3830. Samples (12) reported that N,N′-dichlorourea has higher molar absorptivity than N-chlorourea; no absorbance spectra for tri- or tetrachlorourea are known to have been reported. The pattern of ∆Cl and ∆A240 at pH ) 7.5 indicates that N-chlorourea is unstable in the presence of free chlorine, but these data do not provide clear evidence of the products that form. To further examine the hypothesis that Cl2 is responsible for the first N-chlorination of urea, an experiment was performed in which the N-chlorourea yield was measured after exposure to chlorine for varying concentrations of chloride ion at fixed pH. From this experiment, it was observed that the N-chorourea yield increased as Clconcentration increased (see Figure SI-4). The data presented in Figures 4 and 5 and Figure SI-4 illustrate that pH and chloride ion concentration both affect the rate of chlorine substitution of the urea molecule. Both parameters influence the distribution of free chlorine among its various forms. In particular, as the concentrations of protons (H+) and chloride ion increase, the fraction of free chlorine present as molecular chlorine (Cl2) increases. The first step in the reaction between chlorine and amines is likely to be electrophilic attack of N by chlorine (16). Since Cl2 is a stronger electrophile than HOCl, it is reasonable to expect that the rate of chlorine substitution by Cl2 would be greater than that attributable to HOCl. Figures 4 and 5 (as well as Figure SI-8) suggest that the first chlorine substitution on the urea molecule involves Cl2, not HOCl. Figure 5 (pH 7.5) indicates that once N-chlorourea is formed, it reacts relatively quickly with free chlorine (including HOCl and OCl-) to form other compounds. Moreover, the relatively rapid decay of the A240 signal in Figure 5 suggests that (some of) the products of these reactions are not strong absorbers at λ ) 240 nm. This behavior is consistent with a reaction mechanism wherein the first N-chlorination reaction is relatively slow
The data in Figure 6 were further analyzed for purposes of defining a rate constant for this first step in the reaction between free chlorine and urea. This reaction was assumed to take the form of: KU-Cl
CO(NH2)2 + Cl2 98 H2NCONHCl + Cl2
(1)
where KU-Cl2 ) second-order rate constant for reaction between urea and Cl2 (M-1 · s-1). Expression 1 represents an elementary, bimolecular reaction between urea and Cl2. By extension, this implies a kinetic expression of: d[N-chlorourea] ) KU-Cl2[urea][Cl2] dt FIGURE 6. Measured and simulated values of A245 as a function of time for chlorination of urea based on the following initial conditions: [urea] ) 4 × 10-4 M; [H+] ) 0.01 M; [Cl-] ) 0.01 M; free chlorine ) 10-4 M; T ) 23 °C. (i.e., rate-limiting), but subsequent N-chlorination and hydrolysis steps that lead to NCl3 formation are rapid. The first chlorination reaction would be expected to reduce proton affinity of N in urea (i.e., make the molecule more acidic). By extension, this would lead to an increase in the rate of subsequent N-chlorination steps, presumably until the molecule reached a condition of complete chlorine substitution (i.e., tetrachlorourea). Indeed, this behavior is consistent with the observations of Ricci and Rosi (17, 18). Additional experimental evidence of the differences in the rates of chlorination for urea and N-chlorourea may be found in the Supporting Information (see Figure SI-8). Collectively, the results of these experiments provide general support for the reaction mechanisms that have been proposed by Samples (12) and Li and Blatchley (3). However, the evidence presented above indicates that the first (and rate-limiting) step in this reaction depends on molecular chlorine (Cl2). It is also evident that under conditions that are present in most recreational water settings (i.e., nearneutral to slightly alkaline pH, and relatively low chloride concentration), urea will represent a reservoir of reduced-N. Although it reacts relatively slowly to form intermediates (e.g., NCl3), as well as stable (e.g., NO3-) reaction products, the steady input of urea to pools dictates that it represents an important precursor to formation of these compounds. Stopped-Flow Experiment. The experiment with the stopped-flow device was based on absorbance measurements at λ ) 245 nm. The experiment was conducted at pH ) 2.0 and at a chloride concentration of 0.01 M; under these conditions, free chlorine was present as Cl2 and HOCl, with a negligible concentration of OCl-. In addition, the initial condition for this reaction was a urea concentration of 4.0 × 10-4 M, and a free chlorine concentration of 1.0 × 10-4 M, thereby yielding Cl:P ) 0.25. Figure 6 illustrates the time-course trace of A245 for the reacting solution. A245 signals reported in this figure were zeroed against the solution at t ) 0. For the conditions of the experiment illustrated in Figure 6, equilibrium calculations indicate that 20.6% of the free chlorine was present in the form of Cl2, thereby yielding a molecular chlorine concentration that was orders of magnitude higher than is normally seen in pools. The reaction appears to have reached completion in less than 2 s. It is presumed that this rapid reaction rate was attributable to the relatively high concentration of Cl2. These data support the hypothesis that molecular chlorine is responsible for the first step in the reaction sequence between free chlorine and urea; namely, N-chlorination.
(2)
KU-Cl2 was estimated based on the initial rate of Nchlorourea formation. This approach was used to minimize the effects of absorbance contributions by other reaction products, as well as the effects of decreases in the concentrations of reactants. Algebraic rearrangement of eq 2 in finite-difference form allowed estimation of KU-Cl2 KU-Cl2 ≈
∆[N-chlorourea]/∆t [urea][Cl2]
(3)
The values of [urea] and [Cl2] used in this expression were the initial concentrations for the experiment. Measurements of A245 were converted to [N-chlorourea] based on application of Beer’s law: [N-chlorourea] )
A245 εN-chlorourea,245l
(4)
The numerator in expression 3 was estimated by regression of the [N-chlorourea] measurements for the first 0.10 s of the experiment (see Figure SI-5). On the basis of this approach, the rate constant for reaction 1 was estimated as 2.35 × 104 M-1 · s-1. As a check of the validity of this approach, a finitedifference model was developed to simulate the timedependent concentrations of urea, N-chlorourea, HOCl, and Cl2 in this experiment: [N-chlorourea]i ) [N-chlorourea]i-1 + (KU-Cl2[Cl2]i-1[urea]i-1)
(5)
[free chlorine]i ) [free chlorine]i-1 - ∆[N-chlorourea]i (6) [Cl2]i ) [Cl2]i-1RCl2
(7)
[HOCl]i ) [free chlorine]i - [Cl2]i
(8)
where the subscript i refers to the value of a parameter at time-step i. On the basis of the known initial concentrations of N-chlorourea, urea, Cl2, and HOCl, and the measured values of their respective molar absorptivities, it was possible to simulate the time-dependent trace of A245: A245,i ) l{(εN-chlorourea,245[N-chlorourea]i) + (εurea,245[urea]i) + (εCl2,245[Cl2]i) + (εHOCl,245[HOCl]i)}
(9)
where A245,i refers to absorbance (λ ) 245 nm) at time i (see Figure 6). Figures SI-6 and SI-7 illustrate the methods used to estimate εCl2,245 and εHOCl,245. The simulation of the time-dependent A245 trace based on the initial reaction rate and application of Beer’s law for VOL. 44, NO. 22, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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multicomponent mixtures yielded a slight underestimate of the A245 signal for time less than approximately 1.0 s, and slightly overestimated the A245 signal for later times. The causes of the deviations between the measured and predicted values are not known. However, likely causes were the formation of (absorbing) reaction products that were not included in the model, or error in the initial rate measurement (see Figure SI-5). Nonetheless, the simulation of the timedependent A245 signal presented in Figure 6 appears to account for the majority of the measured dynamics of the system. Overall Reaction Mechanism. By combining the results of his measurements with conjecture regarding reaction progress, Samples (12) hypothesized an overall reaction mechanism for chlorination of urea. The Samples mechanism accounts for much of the observed behavior of this reaction; however, it suggests that the first step in the sequence was N-chlorination by HOCl. The data presented above indicate that this first step in the process, which is also the overall rate-limiting step in this reaction, appears to require molecular chlorine (Cl2) to proceed. On the basis of this information, the hypothesized mechanism of Samples (12) is modified as follows: CO(NH2)2 + Cl2 f H2NCONHCl + Cl-
(10)
H2NCONHCl + HOCl f CO(NHCl)2 + H2O
(11)
CO(NHCl)2 + HOCl f Cl2NCONHCl + H2O
(12)
Cl2NCONHCl + HOCl f CO(NCl2)2 + H2O
(13)
CO(NCl2)2 + HOCl f H+ + Cl- + CO2 + NCl3 + NCl (14) NCl + OH- f NOH + Cl-
(15)
2NOH f H2N2O2
(16)
H2N2O2 f N2O + H2O
(17)
+ NCl3 + HOCl + 2H2O f NO3 + 4Cl + 5H
(18)
Note that the unstable intermediate NOH has also appeared in other mechanisms that address the behavior of reduced-N and free chlorine (15). The involvement of Cl2 in the first (rate-limiting) step of this sequence is important with respect to swimming pool applications. In most pools, only a small fraction of free chlorine will be present as Cl2. As such, urea chlorination will proceed slowly, as suggested by the data reported herein. Some pools employ so-called “salt generation” systems wherein chloride concentration is raised to roughly 10% of the chloride concentration that is typical of seawater, thereby allowing in situ generation of free chlorine by electrolysis. The mechanism described above implies that these conditions will promote chlorine substitution of urea, thereby increasing the rates of NCl3 formation and urea decay. In any case, urea represents an important precursor to NCl3 formation in pools. Education of swimmers on this behavior may encourage reductions in the rates of urea introduction, as well as the rates at which other precursors are introduced. By extension, this may lead to improvements in pool water and air quality.
Acknowledgments The authors received support for this manuscript from the Research Foundation for Health and Environmental Effects
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(RFHEE). RFHEE was not involved in the design, collection, management, analysis, or interpretation of the data; or in the preparation or approval of the manuscript. The findings and conclusions in this manuscript are those of the authors and do not necessarily represent the views of RFHEE.
Supporting Information Available Additional information as noted in the text. This material is available free of charge via the Internet at http://pubs.acs.org.
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