- Email: [email protected]
Kinetics of cyanogen chloride destruction by chemical reduction methods
ARTICLE IN PRESS
Water Research 39 (2005) 2114–2124 www.elsevier.com/locate/watres
Kinetics of cyanogen chloride destruction by chemical reduction methods Chii Shang, Yinan Qi, Li Xie, Wei Liu, Xin Yang Department of Civil Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Kowloon, Hong Kong Received 28 October 2004; received in revised form 25 February 2005; accepted 6 March 2005 Available online 11 May 2005
Abstract In this study, membrane introduction mass spectrometry (MIMS) was applied to evaluate the kinetics of cyanogen chloride (ClCN) destruction by chemical reduction methods, using thiosulfate, sulfite, metabisulfite, ferrous ions and zero-valent iron at various concentrations and pH. The ClCN destruction followed second-order reaction kinetics in all cases of using sulfur compounds, though the second-order rate constants varied substantially from approximately 0.3–25.7 M1 s1 under different experimental conditions. The destruction of ClCN was primarily attributable to the chemical reduction pathway. Hydroxide-assisted ClCN hydrolysis was only significant at pH 9 and also when the observed reduction rate was relatively slow. The second-order rate constants achieved by sulfur(IV) compounds in the form of sulfite were found to be higher than those obtained with thiosulfate and S(IV) compounds in the form of bisulfite. Ferrous ions and zero-valent iron demonstrated slow or no ClCN reduction up to dosages of 1000 mg L1 and 100 g L1, respectively. These findings suggest that applying moderately high dosages of S(IV) compounds under neutral or alkali conditions with sufficient contact time is required for wastewater ClCN destruction. In addition, ClCN losses during long-term preservation with excess reducing sulfur compounds prior to analysis can be substantial and should be avoided. r 2005 Elsevier Ltd. All rights reserved. Keywords: Cyanogen chloride; Chemical reduction; Dechlorination; MIMS; Reducing sulfur compounds
1. Introduction Cyanogen chloride (ClCN), with molecular structure of Cl–CN, is colorless, volatile, and slightly soluble at ambient temperature (25 1C). Similar to cyanide (CN), ClCN is highly toxic to mammal through causing respiratory failure and blocking cellular energy metabolism (American Conference of Governmental Industrial Hygienists, 1991; Lewis, 1996). It is also highly irritate to eyes and mucous membrane therefore has been used as a Corresponding author. Tel.: +852 2358 7885; fax: +852 2358 1534. E-mail address: [email protected] (C. Shang).
chemical warfare agent in World War II (Robinson, 1967). ClCN is very toxic to aquatic life. The 24-h and 48-h LC50s, which are the doses that kill 50% of a population after the specific exposure time, to Daphnia magna (5-day old) were 86 and 65 mg L1, respectively (Kononen, 1988). ClCN is currently unregulated in the United States, but the USEPA currently lists ClCN in the Information Collection Rule (ICR) and considers it as a tentative candidate for regulation (US Environmental Protection Agency (USEPA), 1996). ClCN production results from chlorinating or chloraminating water that contains nitrogen compounds (e.g., amino acids, nucleic acids, and humic acids) (Ohya and Kanno, 1985, 1989; Hirose et al., 1988; Shang et al.,
0043-1354/$ - see front matter r 2005 Elsevier Ltd. All rights reserved. doi:10.1016/j.watres.2005.03.031
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2000). Reacting monochloramine with formaldehyde also leads to the formation of ClCN (Pederson et al., 1999). These precursors have been found and identified at trace quantities in wastewater operations. ClCN has been reported in chloraminated drinking water and wastewater at concentrations of several to a few dozen micrograms per liter (Krasner et al., 1989; Zhang et al., 2004). ClCN can also be an intermediate produced from the first step of alkaline chlorination, which is the most commonly applied technique in treating cyanide-bearing wastewater from metal finishing and metal plating industry (Woodard, 2001). The produced ClCN needs to be rapidly converted to cyanate ions (CNO) by the hydroxide alkalinity. Cyanate ions are roughly a thousand times less toxic than ClCN. Complete oxidation from cyanate ions to nitrogen gas by further chlorination at lower pH (o7.6) may be required. ClCN can be naturally decomposed through hydrolysis, yielding cyanate ions, hydronium ions, and chloride ions as the final products (Pedersen and Marin˜as, 2001). However, the reaction is hydroxideassisted and the half-life of ClCN due to its hydrolysis is on the order of hours at solution pH lower than 10 and at room temperature (Kononen, 1988; Pedersen and Marin˜as, 2001). As such, discharging ClCN containing effluents directly into a receiving water body may impose a threat to downstream aquatic life. In many wastewater treatment facilities that use chlorine to control effluent microbial quality, dechlorination is utilized to eliminate residual chlorine to reduce effluent toxicity. Dechlorination can be achieved by activated carbon, ferrous sulfate, hydrogen peroxide, thiosulfate, and S(IV) compounds (White, 1999). Among the listed substances, S(IV) compounds are the most commonly used reducing reagents in wastewater operations while thiosulfate is frequently used in laboratory settings. Most residual chlorine is rapidly quenched (eliminated) when S(IV) compounds are well mixed with water. It has been demonstrated that the reduction of effluent toxicity is also associated with destruction of disinfection by-products (DBPs). Yields of trihalomethanes (THMs) such as chloroform and bromodichloromethane can be reduced by sulfite (Helz et al., 1985), and reduction of total organic halogen (TOX) and mutagenicity has been reported by achieving complete dechlorination with sulfite (Morlay et al., 1991). In this study, therefore, the kinetics of ClCN destruction by chemical reduction methods and the corresponding rate constants with variations in dechlorination reagents and pH were evaluated. These dechlorination reagents includes thiosulfate, sulfite, metabisulfite, ferrous ions and zero-valent iron. ClCN destruction by S(IV) compounds at neutral or alkali pH could be a good method for reducing toxicity in
2115
chlorinated wastewater effluent and for neutralizing ClCN as a chemical warfare agent.
2. Materials and methods 2.1. Solution preparation In all cases, solutions were prepared from reagentgrade chemicals or stock solutions. Dilution to target aqueous-phase concentrations was accomplished with deionized, double-distilled water. The stock free chlorine (HOCl) solution was prepared from 4% sodium hypochlorite (NaOCl) (Aldrich), diluted to approximately 4000 mg L1 as Cl2, and stored in an aluminum foilcovered glass-stoppered flask. The stock solution was periodically standardized by DPD/FAS titration (Standard Methods for the Examination of Waste and Wastewater, 1998). The glycine solution, used as a precursor to yield ClCN from chlorination, was prepared from a stock solution (100 mg L1 as N) obtained from a reagent-grade chemical (Sigma). Sodium thiosulfate (Na2S2O3) (Fluka), sodium sulfite (Na2SO3) (Nacalai Tesque), sodium metabisulfite (Na2S2O5) (Sigma), and ferrous chloride (FeCl2) (Riedel-deHae¨n) were freshly dissolved into dilution water to make stock solutions of predetermined concentrations. Precisely weighted zero-valent iron (Fe0) powders (Connelly-GPM ETICC-1004) were directly applied to a test solution at predetermined solid-to-liquid ratios (concentrations). Due to the prohibition on importing ClCN to Hong Kong and to the short shelf-life of ClCN, standards of ClCN for calibration curve establishment were freshly, in situ developed by 30-min chlorination of glycine solution (0.5 mg L1 as N) at a free chlorine dosage of 5 mg L1 (as Cl2) and phosphate-buffered pH of 7. This protocol had demonstrated to yield consistent aqueous ClCN concentrations of 900 mg L1 while leaving no residual chlorine in the solution (Shang et al., 2000). Thereafter, dilutions were made to give standards at 900, 100, 10, and 2 mg L1. 2.2. Analytical methods The membrane introduction mass spectrometry (MIMS) method was applied for on-line, rapid aqueous ClCN monitoring and analysis. MIMS employs a membrane interface to introduce an aqueous solution, without any sample work-up, directly to a mass spectrometer by allowing the ClCN to permeate continuously through the membrane while water and ionic species are rejected (Shang et al., 2000; Pederson et al., 1999; Yang and Shang, 2004, 2005). The detailed description of the MIMS method for ClCN quantification can be found in Yang and Shang (2005). In brief, it
ARTICLE IN PRESS 2116
C. Shang et al. / Water Research 39 (2005) 2114–2124
was based on a modification of an HP 5892 bench-top GC/MS (Hewlett-Packard) containing an HP 5972A mass selective detector (MSD) using electron ionization (70 eV) and the selected ion monitoring (SIM) mode. ClCN concentrations were directly determined by comparison of the ion abundance measurements at m/z 61 amu (35ClCN+) with those developed from a series of standard solutions, as the isotopic ion (37ClCN+) at m/z 63 amu may interfere with the molecular ion of an intermediate, N-chloromethanimine (CH2N35Cl+) (Yang and Shang, 2005). The limit of detection of ClCN was 1 mg L1. 2.3. Experimental procedures Most experimental runs were conducted in a 0.5-L well-mixed (using a magnetic stirrer and a stirring bar), parafilm-sealed glass reactor with a 500 mL aqueous glycine solution (0.5 mg L1 as N), phosphate-buffered to pH 5, 6.9, or 9 (0.01 M) at room temperature. The solution pH was periodically checked and no change of the pH was observed in all cases. The glycine solution was first pumped to the MIMS system for 5 min, at which time NaOCl (5 mg L1 as Cl2) was added (the first arrow in all figures) for 30 min to synthesize ClCN. The concentration of ClCN increased with time until it reached a stable, maximum ClCN yield approximately 20 min after the chlorine addition. Thereafter, the concentration of ClCN remained relatively constant until a specific dechlorination reagent was introduced to the testing solution at 35 min (the second arrow in all figures). At the time of introducing the dechlorination reagent, no residual chlorine remained in the solutions. The introduction of the dechlorination reagent required that the parafilm seal be broken and the reactor be immediately resealed. The ClCN concentration was continuously monitored by MIMS until the test run terminated at 70 min. A few test runs were repeated in a 1-L well-mixed, head-space-free, Teflon bag reactor to confirm that the small head space in the glass reactor and breakage of the parafilm seal unaffected the test results. In the cases of using Fe0 powders, a membrane filter with a pore size of 0.45 mm was installed in the sampling line to prevent damage or fouling of the membrane interface from the insoluble iron.
3. Results and discussion ClCN destruction and its kinetics by the three sulfur compounds at pH 5, 6.9, and 9 were first evaluated. Dechlorination processes using reducing sulfur compounds have been well developed to eliminate chlorine residuals and can be easily applied with commercially available packages. Concentrations of the reducing sulfur compounds used in this study were much higher
than those in the practical range to establish the kinetic rate constants. ClCN destruction with the sulfur compounds was assumed to follow the common dechlorination principle that the chlorine atom of ClCN is dechlorinated and reduced to chloride ions (Cl), as S2 O2 3 (or other sulfur compounds) is oxidized to SO2 : 4 k
2 2 2 S2 O2 3 ðSO3 or S2 O5 Þ þ ClCN ! SO4 þ Cl þ other products;
ð1Þ
where k is the second-order reaction rate constant, by assuming a second-order bimolecular reaction between ClCN and the sulfur compound employed in Eq. (1). The rate of ClCN destruction can be then expressed as d½ClCN ¼ k½ClCN ½reducing sulfur compound dt
(2)
where [i] denotes the molar concentration of compound i and t is time. The destruction of ClCN through the reaction was further characterized by assuming pseudofirst-order decay at high concentrations of the sulfur compounds. As such, the rate of ClCN destruction can be expressed as d½ClCN ¼ kobs ½ClCN ; dt
(3)
where kobs is the observed pseudo-first-order rate constant and is defined as the product of the secondorder rate constant (k) and the initial concentration of the reducing sulfur compound. The kobs therefore can be obtained by the slope of the Ln[ClCN] versus t plot. ClCN is also known to hydrolyze in alkali solution according to the general base-catalyzed reaction (Pedersen and Marin˜as, 2001): kh
ClCN þ H2 O ! HOCN þ Hþ þ Cl
(4)
in which kh is the overall hydrolysis rate constant defined as B kh ¼ koh þ kOH h ½OH þ kh ½B ,
koh ,
kOH h ,
(5)
kBh
and are the water, hydroxide ion, where and generic base-B-assisted rate constants. The koh ranging from 3.29 107 to 2.58 106 s1 and the kOH ranging from 3.5 to 13 M1 s1 have been reported h in the literature (Pedersen and Marin˜as, 2001). Thus, the kobs can be rewritten as kobs ¼ kred þ kh ,
(6)
where kred is the pseudo-first-order rate constant contributed only from the chemical reduction reaction. Figs. 1–3 present the results of the experimental runs using sodium thiosulfate, sodium sulfite, and sodium metabisulfite, respectively. Table 1 lists all calculated
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124 900
pH 5.0
2117
Ion 61 (CN35Cl+•)
800
Conc. of CNCl (µg/L)
700 600 500
1.14 mM Na2S2O3
400 300 200 100
11.4 mM Na2S2O3 0 10.00
20.00
30.00
(a)
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500
1.14 mM Na2S2O3
400 300 200 100
11.4 mM Na2S2O3
0 10.00
20.00
30.00
(b)
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 40
pH 9.0
Conc. of CNCl (µg/L)
35 30 25 20 15
1.14 mM Na2S2O3
10 5
11.4 mM Na2S2O3
0 10.00
20.00
30.00
(c)
40.00
50.00
60.00
70.00
Time (minute)
Fig. 1. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium thiosulfate. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O3.
rate constants and R2 corresponding to the experimental results obtained with variations in species and concentrations of the reducing sulfur compounds and varying
pH levels. As listed in Table 1, in all cases, an R2 of 0.946 or higher is obtained, indicating that the pseudo-firstorder assumption is valid.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2118 900
Ion 61 (CN35Cl+•)
pH 5.0
800
Conc. of CNCl (µg/L)
700 600 500
0.633 mM Na2SO3
400 300 200
1.266 mM Na2SO3
100 0 10.00
20.00
(a)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500 400 300 200
0.633 mM Na2SO3 1.266 mM Na2SO3
100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
70.00
Time (minute)
40
Ion 61 (CN35Cl+•)
pH 9.0
Conc. of CNCl (µg/L)
35 30 25 20 15 10
0.378 mM Na2SO3
5
1.266 mM Na2SO3 0 10.00
(c)
20.00
30.00
40.00
50.00
60.00
70.00
Time (minute)
Fig. 2. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium sulfite. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2SO3.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124 900
pH 5.0
2119
Ion 61 (CN35Cl+•)
800
Conc. of CNCl (µg/L)
700 600 500
0.316 mM Na2S2O5
400
1.266 mM Na2S2O5
300 200 100 0 10.00
20.00
(a)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500 400 300 200
0.316 mM Na2S2O5 1.266 mM Na2S2O5
100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 40
pH 9.0 0 mM Na2S2O5
Conc. of CNCl (µg/L)
35 30 25 20 15 10 5
1.266 mM Na2S2O5
0.316 mM Na2S2O5
0 10.00
(c)
20.00
30.00
40.00
50.00
60.00
70.00
Time (minute)
Fig. 3. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium metabisulfite. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O5.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2120
Table 1 Pseudo-first-order rate constant (kobs), second-order rate constant (k), and R2 for regression of ClCN destruction with the three reducing sulfur compounds at pH 5, 6.9, and 9
(a) Na2S2O3
pH
Initial conc. (mM)
Calculated Kobs (min1)
R2
5.0
1.14 11.4 1.14 11.4 1.14 11.4
0.04 0.42 0.04 0.41 0.06 0.42
0.998 0.993 0.999 0.986 0.999 0.991
0.61 0.61 0.61 0.61 0.89 0.61
0.633 1.266 0.633 1.266 0.380 1.266
0.03 0.06 0.14 0.26 0.30 0.87
0.999 0.999 0.999 0.991 0.989 0.986
0.81 0.79 3.66 3.36 13.1 11.4
0.316 1.266 0.316 1.266 0.316 1.266
0.01 0.02 0.11 0.34 0.52 1.95
0.994 0.967 0.994 1.000 0.957 0.946
0.62 0.31 5.60 5.26 27.2 25.7
6.9 9.0 (b) Na2SO3
5.0 6.9 9.0
(c) Na2S2O5
5.0 6.9 9.0
3.1. Dechlorination with sodium thiosulfate Fig. 1(a–c) present the results of ClCN destruction by sodium thiosulfate at pH 5, 6.9, and 9, respectively. The thiosulfate additions varied at 1.14 and 11.4 mM. Generally, although there was a reaction, the carbonbonded chlorine reacted slowly with thiosulfate. Only thiosulfate addition at a high, unpractical concentration (11.4 mM) completely destroyed ClCN in 15 min at pH 6.9. Similar trends were found when the experiments were conducted at buffered pH of 5 and 9 while other conditions remained unchanged. It should be noted that, at pH 9, a sudden drop in the concentration of ClCN was observed right after the addition of the thiosulfate (as shown in Fig. 1(c)). The sudden drop was initially postulated to be due to the breakage of the seal for injecting the thiosulfate. However, the hypothesis was denied by the later confirmation tests using a head-space free, sealed Teflon bag reactor and syringe-injection of the sulfur compound where the sudden drop remained. However, we are unable to explain the observation and such a sharp decrease of ClCN was excluded from the data analysis, considering that the later stage is the key stage for ClCN destruction. Table 1(a) presents the calculated rate constants of ClCN destruction by thiosulfate. At pH 5 and 6.9, the calculated second-order rate constants were identical (0.61 M1 s1) regardless of the variations of thiosulfate concentrations and pH. The result indicates that the same mechanism (ClCN destruction mainly by thiosulfate) is applied and ClCN hydrolysis can be neglected at
Calculated k (M1 s1)
the pH range. In other words, kobsEkred. Using the 6 1 largest values of the koh and the kOH s and h , 2.58 10 1 1 13 M s , respectively, reported in the literature (Pedersen and Marin˜as, 2001), the calculated kh is at least two orders of magnitude smaller than the kred at pH 7 or less. On the other hand, considerable discrepancy was observed at pH 9 wherein k’s of 0.89 and 0.61 M1 s1 were obtained at initial thiosulfate concentrations of 1.14 and 11.4 mM, respectively. The discrepancy can be explained by the extent of the incorporation of ClCN hydrolysis in the destruction process. The relative significance of ClCN hydrolysis on ClCN destruction depends on the relative time scale of the reactions. At high initial thiosulfate concentration (11.4 mM), the dechlorination reaction was fast and ClCN completely disappeared in less than 10 min so that the reaction rate constant was not affected by the ClCN hydrolysis. However, ClCN hydrolysis greatly enhanced the ClCN destruction at low thiosulfate addition (1.14 mM) and pH 9. The average kh at pH 9 was calculated from subtraction (based on Eq. (6)) to be 0.01 min1. The somewhat higher kh at pH 9 obtained in the current study is due to the use of a phosphate buffer. Phosphate at moderate pH (X6.5) has been reported to give an increase in the rate of ClCN hydrolysis (Edwards et al., 1986). 3.2. Dechlorination with sodium sulfite Similar experimental runs were conducted using sodium sulfite (0.633 and 1.266 mM) (see Fig. 2). After
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
dissolution, Na2SO3 is expected to yield an equilibrium 2 mixture of bisulfite (HSO 3 ) and sulfite (SO3 ) in aqueous solution according to its acidity constant (Benjamin, 2001): Ka
2 þ HSO 3 2 H þ SO3 ;
pK a ¼ 7:3.
(7)
As such, the reduction of ClCN is attributable to the dechlorination reactions by bisulfite, sulfite or both based on the solution pH. ClCN hydrolysis may also contribute to the overall ClCN reduction. As shown in Fig. 2, although trends that ClCN destruction rates increased with increases in sodium sulfite dosage and pH were observed, the magnitude of the rate increases with increasing pH was much greater than those obtained with sodium thiosulfate additions and the increase also occurred even at pH changes from 5 to 6.9, at which the ClCN hydrolysis could be neglected. The rate constant calculation (see Table 1(b)) followed the same manner as described above and the second-order rate constants at specific pH agreed reasonably well when comparing the values obtained at different initial dosages. The significant rate increases (from approximately 0.8 to 12 M1 s1) with increasing pH were further demonstrated in the rate constant calculation, indicating that ClCN destruction by sodium sulfite is much more sensitive to the solution pH. To shed some light on the causes, calculation on the overall rate constant expressed as the sum of the rates of reactions with sulfite and bisulfite (modified from Croue and Reckhow, 1989) was adopted: 2 kobs ¼ k1 ½HSO 3 þ k2 ½SO3 þ kh ,
(8)
where the kh can be neglected since kh5kobs and k1 and k2 can be obtained as 0.80 and 12.3 M1 s1 at pH 5 and 9, respectively, when the contribution of one or the other is negligible due to its extremely low concentration. The substantial difference in k1 and k2 is expected since sulfite is a much stronger nucleophile compared to bisulfite (Croue and Reckhow, 1989). Correspondingly, the calculated overall second-order rate constant (k) at pH 6.9 is 4.06 M1 s1, which is not far from the experimental data shown in Table 1(b). These findings suggest that the speciation of S(IV) compounds, rather than ClCN hydrolysis, contributes to the variation in the destruction rates. Compared with ClCN destruction by thiosulfate (see Fig. 1 and Table 1), at pH 6.9 and 9, the dose requirement of sodium sulfite to achieve the same degree of ClCN reduction was much smaller and the secondorder rate constants of ClCN destruction by sodium sulfite were much higher than those of thiosulfate. However, the difference was not noticeable at pH 5, where mainly bisulfite was predominant, indicating that bisulfite is as weak as thiosulfate with regard to ClCN dechlorination. These findings only partially agree with
2121
the literature (White, 1999) that S(IV) compounds are better than thiosulfate for dechlorination. It suggests that only S(IV) compounds in the form of sulfite (SO2 3 ) are effective for destruction of ClCN (presumably other DBPs as well) and controlling the pH at neutral or alkali conditions to yield a higher fraction of sulfite is essential. In addition, providing sufficient contact time is also required. To evaluate the effect of residual sulfite/bisulfite on dissolved oxygen (DO) levels, DO was measured at a sodium sulfite dosage of 1.266 mM and pH 7; and DO reduction from 9 to 3.4 mg L1 was observed at the end of the test run. 3.3. Dechlorination with sodium metabisulfite Sodium metabisulfite (0.316 and 1.266 mM) was tested under similar experimental conditions. As shown in Fig. 3 and Table 1(c), similar to the results of using sodium sulfite, large enhancements in ClCN reduction were observed with increases in metabisulfite additions and pH. Again, the second-order rate constants of the S(IV) compound at specific pH agree reasonably well at different initial dosages. Nevertheless, at the same dosage, ClCN reduction by sodium metabisulfite was notably faster (close to double) than that achieved by sodium sulfite at pH 6.9 and 9. This finding can be explained from the equilibrium between bisulfite and metabisulfite (Yiin et al., 1987): keq
2 2HSO 3 2 S2 O5 þ H2 O;
keq ¼ 6:5 102 M1
at 25 C:
ð9Þ
At the metabisulfite concentrations used (0.316 and 1.266 mM), the equilibrium should produce mainly sulfite/bisulfite in stoichiometric quantities with negligible concentrations of metabisulfite and there should be an equilibrium (Eq. (7)) between sulfite and bisulfite according to the solution pH. To directly obtain the hydrolysis rate constant, kh, at pH 9, a side experiment was conducted without adding any reducing sulfur compound. As shown in Fig. 3, after 35 min, approximately 25% of the produced ClCN was lost from the hydrolysis. By assuming a first-order reaction, the kh was estimated to be 0.01 min1, which was in agreement with that previously estimated from subtraction. Interference at the second highest peak of ClCN (corresponding to CN37Cl+ at m/z 63 amu) concurrent with the presence of a considerable abundance of ions at m/z 65 amu was also observed in all cases. The abundance at m/z 65 amu was attributed to the formation of an intermediate, N-chloromethanimine (CH2N37Cl) (Yang and Shang, 2005). As shown in Fig. 4 as an example of using Na2S2O5, a relatively much sharper decrease in the abundance of m/z 65 amu
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2122
Abundance
Ion 61.00 (60.70 to 61.70): 200B.D CN35Cl+•
5500 5000 4500 4000 3500 3000 2500 2000 1500 1000 500 10.00
20.00
(a)
30.00 40.00 Time-->
50.00
60.00
70.00
Abundance
Ion 65.00 (64.70 to 65.70): 200B.D 350 300 250 200 150 100
CH2N37Cl+•
10.00
20.00
30.00
(b)
40.00
50.00
60.00
70.00
Time-->
Fig. 4. On-line selected ion monitoring of chlorinating/dechlorinating glycine solution (0.5 mg L1 as N) at pH 6.9. Ion chromatogram presented are (a) m/z 61 amu (CN35Cl+) for ClCN and (b) m/z 65 amu (CH2N37Cl+) for CH2NCl. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O5 (1.266 mM).
was observed immediately after reducing sulfur compound additions. This finding indicates that this nitrogen-bonded chlorine reacts with reducing sulfur compounds much more quickly than the carbon-bonded chlorine in ClCN. Therefore, the nitrogen–chlorine bond is more vulnerable to breakage than the carbon– chlorine bond during the dechlorination process. Bond dissociation energies of 333.979.6 and 397 729 kJ mol1 for the nitrogen–chlorine bond and the carbon–chlorine bond, respectively, at 25 1C are reported by Kerr (2000). In addition, the half-life of N-chloromethanimine (approximately 1 min) obtained from this study during dechlorination was consistent with that of the organic chloramines (on the order of sub-minutes to a few minutes) reported in the literature (Helz and Nweke, 1995; Jensen and Helz, 1998; Jameel and Helz, 1999) at similar S(IV) dosages. Therefore, when carbon-bonded chlorine is required to be effectively destructed by a common dechlorination process using the reducing sulfur compounds, the condition should be likely able to dechlorinate organic chloramines. Further investigation is needed to verify this assertion.
Shang, 2005; Gordon et al., 2002; Orth and Gillham, 1996). Fe2+ and Fe0 were postulated to be capable of reducing ClCN and experimental trials have been conducted. The trials were managed in the same manner as the previous cases except Fe2+ at 50 and 1000 mg L1 and Fe0 at 10, 50, and 100 g L1 were used at pH 6.9. Figs. 5a and b present the resulting ClCN concentrations treated by Fe2+ and Fe0, respectively. As shown, after 35 min, no ClCN could be reduced by up to 1000 mg L1 of Fe2+, while only approximately 50% of ClCN could be reduced by 100 g L1 Fe0. For comparison, 15 mg L1Fe2+ and 50 g L1 of the same Fe0 powders were found to reduce bromate from 50 to 10 mg L1 within 15 min (Siddiqui et al., 1994) and from 100 to 10 mg L1 within 10 min (Xie and Shang, submitted), respectively. ClCN destruction was found to be more feasible from using S(IV) compounds than Fe2+ in this study, which was contradictory to the opposite tendency of bromate destruction from S(IV) compounds and Fe2+ documented in the literature (Gordon et al., 2002). No conclusive explanation for this can be offered here.
3.4. Dechlorination with ferrous ions and zero-valent iron
4. Conclusions
Both ferrous ions (Fe2+) and zero-valent iron (Fe0) are powerful reducing agents and can effectively remove bromate (BrO 3 ) and chlorinated compounds (Xie and
Reducing sulfur compounds may be effective in eliminating ClCN in aqueous solution. In all cases tested, the ClCN destruction was primarily attributable
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2123
Ion 61 (CN35Cl+•) 900
1000 mg/L Fe2+
Fe2+ 50 mg/L Fe2+
Conc. of CNCl (µg/L)
800 700 600 500 400 300 200 100 0 10.00
20.00
30.00
40.00
50.00
60.00
Time (minute)
(a)
Ion 61 (CN35Cl+•) 900
Fe0
10 g/L Fe0
Conc. of CNCl (µg/L)
800
50 g/L Fe0
700 600
100 g/L Fe0
500 400 300 200 100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
Time (minute)
Fig. 5. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating by ferrous ions or zero-valent iron at pH 6.9. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Fe2+ (or Fe0).
to the chemical reduction pathway and followed secondorder reaction kinetics, though the second-order rate constants varied substantially from approximately 0.3 to 25.7 M1 s1 under different experimental conditions. Hydroxide-assisted ClCN hydrolysis was only significant at pH 9 and also when the observed chemical reduction rate was relatively slow. Among the tested compounds and pH, the S(IV) compound in the form of sulfite (SO2 3 ) displays the best destruction rate. However, the observed slower ClCN reduction rates, compared with those of organic chloramines, suggest that ClCN is likely to contribute to effluent toxicity in current wastewater chlorination/dechlorination practices. Therefore, applying a moderately higher amount of the S(IV) compounds (either sulfite or metabisulfite) at neutral or alkali pH and/or an effluent holding tank to increase the contact time before final discharge may be the possible solution(s) to achieve complete ClCN destruction. Nevertheless, it should be noted that the dechlorination of free chlorine and chloramines has been
found to increase with decreasing pH (Yiin et al., 1987; Yiin and Margerum, 1988). An elevated pH may slow down the removal of free chlorine and inorganic and organic chloramines, which are expected to be the major fractions of chlorine after wastewater chlorination. Balance should be carefully sought. On the other hand, slow ClCN destruction is expected even with a small dosage of the three sulfur compounds, indicating that the common practice of using reducing sulfur compounds (e.g., Na2S2O3) to quench chlorination reactions before analysis may cause significant ClCN losses (presumably losses of other DBPs as well), if long-term preservation and storage are not properly avoided.
Acknowledgements The authors are grateful to the Research Grants Council, Hong Kong for financial support of the project
ARTICLE IN PRESS 2124
C. Shang et al. / Water Research 39 (2005) 2114–2124
under project no. DAG00/01EG01 and to Professor Blatchley III ER of the Purdue University, School of Civil Engineering for his input in developing the MIMS method and at the early stage of the preliminary trials.
References American Conference of Governmental Industrial Hygienists, 1991. Documentation of the Threshold Limit Values ad Biological Exposure Indices, sixth ed. Cincinnati, OH. Benjamin, M.M., 2001. Water Chemistry. Mc-Graw Hill, New York. Croue, J.P., Reckhow, D.A., 1989. Destruction of chlorination byproducts with sulfite. Environ. Sci. Technol. 23 (11), 1412–1419. Edwards, J.O., Erstfeld, T.E., Ibne-rasa, K.M., Levey, G., Moyer, M., 1986. Reaction rates for nucleophiles with cyanogen chloride: comparison with two other diagonal carbon compounds. Int. J. Chem. Kinetics 18, 165–180. Gordon, G., Gauw, R.D., Emmert, G.L., Walters, B.D., Bubnis, B., 2002. Chemical reduction methods for bromate ion removal. J. Am. Water Works Assoc. 94 (2), 91–98. Helz, G.R., Nweke, A.C., 1995. Incompleteness of wastewater dechlorination. Environ. Sci. Technol. 29 (4), 1018–1022. Helz, G.R., Uhler, A.D., Sugam, R., 1985. Dechlorination and trihalomethane yields. Bull. Environ. Contam. Toxicol. 34, 497–503. Hirose, Y., Maeda, N., Ohya, T., Nojima, K., Kanno, S., 1988. Formation of cyanogen chloride by the reaction of amino acids with hypochlorous acid in the presence of ammonium ion. Chemosphere 17 (5), 865–873. Jameel, R.H., Helz, G.R., 1999. Organic chloramines in disinfected wastewaters: Rates of reduction by sulfite and toxicity. Environ. Toxicol. Chem. 18 (9), 1899–1904. Jensen, J.S., Helz, G.R., 1998. Rates of reduction of Nchlorinated peptides by sulfite: Relevance to incomplete dechlorination of wastewaters. Environ. Sci. Technol. 32 (4), 516–522. Kerr, J.A., 2000. Strengths of chemical bonds. In: Lide, D.R. (Ed.), CRC Handbook of Chemistry and Physics 1999–2000, 81st ed. CRC Press, Boca Raton, Florida. Kononen, D.W., 1988. Acute toxicity of cyanogen chloride to Daphnia magna. Bull. Environ. Contam. Toxicol. 41 (3), 371–377. Krasner, S.W., McGuire, M.J., Jacangelo, J.G., Patania, N.L., Reagan, K.M., Aieta, E.M., 1989. The occurrence of disinfection by-products in US drinking water. J. Am. Water Works Assoc. 81 (8), 41–53. Lewis Sr., R.J., 1996. Sax’s Dangerous Properties of Industrial Materials, ninth ed. Van Nostrand Reinhold, New York, NY. Morlay, C., de Laat, J., Dore´, M., Courtois, Y., Houel, N., Montiel, A., 1991. Effect of an addition of sodium-sulfite on the mutagenicity of chlorinated solutions of aquatic humic substances. Bull. Environ. Contam. Toxicol. 47, 15–22.
Ohya, T., Kanno, S., 1985. Formation of cyanide ion or cyanogen chloride through the cleavage of aromatic rings by nitrous acid or chlorine. VIII. On the reaction of humic acid with hypochlorous acid in the presence of ammonium ion. Chemosphere 14, 1717–1722. Ohya, T., Kanno, S., 1989. Formation of cyanide ion or cyanogen chloride through the cleavage of aromatic rings by nitrous acid or chlorine. XI. On the reaction of purine bases with hypochlorous acid in the presence of ammonium ion. Chemosphere 19 (12), 1835–1842. Orth, W.S., Gillham, R.W., 1996. Dechlorination of trichloroethene in aqueous solution using Fe-0. Environ. Sci. Technol. 30 (1), 66–71. Pedersen III, E.J., Marin˜as, B.J., 2001. The hydroxide-assisted hydrolysis of cyanogen chloride in aqueous solution. Water Res. 35 (3), 643–648. Pederson III, E.J., Urbansky, E.T., Marin˜as, B.J., Margerum, D.W., 1999. Formation of cyanogen chloride from the reaction of monochloramine with formaldehyde. Environ. Sci. Technol. 33 (23), 4239–4249. Robinson, J.P., 1967. Chemical warfare. J. Sci. 3, 33–40. Shang, C., Gong, W-L., Blatchley III, E.R., 2000. Breakpoint chemistry and volatile byproduct formation resulting from chlorination of model organic-N compounds. Environ. Sci. Technol. 34 (9), 1721–1728. Siddiqui, M., Amy, G., Ozekin, K., Zhai, W., Westerhoff, P., 1994. Alternative strategies for removing bromate. J. Am. Water Works Assoc. 86 (10), 81–96. Standard Methods for the Examination of Water and Wastewater, 1998. Standard Methods for the Examination of Water and Wastewater, 20th ed. APHA, AWWA, and WEF, Washington, DC. US Environmental Protection Agency (USEPA), 1996. Fed. Regist. 61, 24353–24388. White, G.C., 1999. Handbook of Chlorination and Alternative Disinfectants, fourth ed. Wiley, New York. Woodard, F., 2001. Industrial Waste Treatment Handbook. Butterworth-Heinemann, Woburn, MA. Xie, L., Shang, C. Bromate reduction by zero valent iron in batch tests. J. Am. Water Works Assoc., submitted for publication. Yang, X., Shang, C., 2004. Chlorination byproduct formation in the presence of humic acid, model nitrogenous organic compounds, ammonia, and bromide. Environ. Sci. Technol. 38 (19), 4995–5001. Yang, X., Shang, C., 2005. Quantification of aqueous cyanogen chloride and cyanogen bromide in environmental samples by MIMS. Water Res., in press. Yiin, B.S., Margerum, D.W., 1988. Kinetics of hydrolysis of the chlorosulfate ion. Inorg. Chem. 27, 1670–1672. Yiin, B.S., Walker, D.M., Margerum, D.W., 1987. Non-metal redox kinetics: general-acid-assisted reactions of chloramine with sulfite and hydrogen sulfite. Inorg. Chem. 26, 3435–3441. Zhang, A., Dzombak, D.A., Luthy, R.G., 2004. Formation of free cyanide and cyanogen chloride from chloramination of publicly owned treatment works secondary effluent: laboratory study with model compounds. Water Environ. Res. 76 (2), 113–120.
Water Research 39 (2005) 2114–2124 www.elsevier.com/locate/watres
Kinetics of cyanogen chloride destruction by chemical reduction methods Chii Shang, Yinan Qi, Li Xie, Wei Liu, Xin Yang Department of Civil Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Kowloon, Hong Kong Received 28 October 2004; received in revised form 25 February 2005; accepted 6 March 2005 Available online 11 May 2005
Abstract In this study, membrane introduction mass spectrometry (MIMS) was applied to evaluate the kinetics of cyanogen chloride (ClCN) destruction by chemical reduction methods, using thiosulfate, sulfite, metabisulfite, ferrous ions and zero-valent iron at various concentrations and pH. The ClCN destruction followed second-order reaction kinetics in all cases of using sulfur compounds, though the second-order rate constants varied substantially from approximately 0.3–25.7 M1 s1 under different experimental conditions. The destruction of ClCN was primarily attributable to the chemical reduction pathway. Hydroxide-assisted ClCN hydrolysis was only significant at pH 9 and also when the observed reduction rate was relatively slow. The second-order rate constants achieved by sulfur(IV) compounds in the form of sulfite were found to be higher than those obtained with thiosulfate and S(IV) compounds in the form of bisulfite. Ferrous ions and zero-valent iron demonstrated slow or no ClCN reduction up to dosages of 1000 mg L1 and 100 g L1, respectively. These findings suggest that applying moderately high dosages of S(IV) compounds under neutral or alkali conditions with sufficient contact time is required for wastewater ClCN destruction. In addition, ClCN losses during long-term preservation with excess reducing sulfur compounds prior to analysis can be substantial and should be avoided. r 2005 Elsevier Ltd. All rights reserved. Keywords: Cyanogen chloride; Chemical reduction; Dechlorination; MIMS; Reducing sulfur compounds
1. Introduction Cyanogen chloride (ClCN), with molecular structure of Cl–CN, is colorless, volatile, and slightly soluble at ambient temperature (25 1C). Similar to cyanide (CN), ClCN is highly toxic to mammal through causing respiratory failure and blocking cellular energy metabolism (American Conference of Governmental Industrial Hygienists, 1991; Lewis, 1996). It is also highly irritate to eyes and mucous membrane therefore has been used as a Corresponding author. Tel.: +852 2358 7885; fax: +852 2358 1534. E-mail address: [email protected] (C. Shang).
chemical warfare agent in World War II (Robinson, 1967). ClCN is very toxic to aquatic life. The 24-h and 48-h LC50s, which are the doses that kill 50% of a population after the specific exposure time, to Daphnia magna (5-day old) were 86 and 65 mg L1, respectively (Kononen, 1988). ClCN is currently unregulated in the United States, but the USEPA currently lists ClCN in the Information Collection Rule (ICR) and considers it as a tentative candidate for regulation (US Environmental Protection Agency (USEPA), 1996). ClCN production results from chlorinating or chloraminating water that contains nitrogen compounds (e.g., amino acids, nucleic acids, and humic acids) (Ohya and Kanno, 1985, 1989; Hirose et al., 1988; Shang et al.,
0043-1354/$ - see front matter r 2005 Elsevier Ltd. All rights reserved. doi:10.1016/j.watres.2005.03.031
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2000). Reacting monochloramine with formaldehyde also leads to the formation of ClCN (Pederson et al., 1999). These precursors have been found and identified at trace quantities in wastewater operations. ClCN has been reported in chloraminated drinking water and wastewater at concentrations of several to a few dozen micrograms per liter (Krasner et al., 1989; Zhang et al., 2004). ClCN can also be an intermediate produced from the first step of alkaline chlorination, which is the most commonly applied technique in treating cyanide-bearing wastewater from metal finishing and metal plating industry (Woodard, 2001). The produced ClCN needs to be rapidly converted to cyanate ions (CNO) by the hydroxide alkalinity. Cyanate ions are roughly a thousand times less toxic than ClCN. Complete oxidation from cyanate ions to nitrogen gas by further chlorination at lower pH (o7.6) may be required. ClCN can be naturally decomposed through hydrolysis, yielding cyanate ions, hydronium ions, and chloride ions as the final products (Pedersen and Marin˜as, 2001). However, the reaction is hydroxideassisted and the half-life of ClCN due to its hydrolysis is on the order of hours at solution pH lower than 10 and at room temperature (Kononen, 1988; Pedersen and Marin˜as, 2001). As such, discharging ClCN containing effluents directly into a receiving water body may impose a threat to downstream aquatic life. In many wastewater treatment facilities that use chlorine to control effluent microbial quality, dechlorination is utilized to eliminate residual chlorine to reduce effluent toxicity. Dechlorination can be achieved by activated carbon, ferrous sulfate, hydrogen peroxide, thiosulfate, and S(IV) compounds (White, 1999). Among the listed substances, S(IV) compounds are the most commonly used reducing reagents in wastewater operations while thiosulfate is frequently used in laboratory settings. Most residual chlorine is rapidly quenched (eliminated) when S(IV) compounds are well mixed with water. It has been demonstrated that the reduction of effluent toxicity is also associated with destruction of disinfection by-products (DBPs). Yields of trihalomethanes (THMs) such as chloroform and bromodichloromethane can be reduced by sulfite (Helz et al., 1985), and reduction of total organic halogen (TOX) and mutagenicity has been reported by achieving complete dechlorination with sulfite (Morlay et al., 1991). In this study, therefore, the kinetics of ClCN destruction by chemical reduction methods and the corresponding rate constants with variations in dechlorination reagents and pH were evaluated. These dechlorination reagents includes thiosulfate, sulfite, metabisulfite, ferrous ions and zero-valent iron. ClCN destruction by S(IV) compounds at neutral or alkali pH could be a good method for reducing toxicity in
2115
chlorinated wastewater effluent and for neutralizing ClCN as a chemical warfare agent.
2. Materials and methods 2.1. Solution preparation In all cases, solutions were prepared from reagentgrade chemicals or stock solutions. Dilution to target aqueous-phase concentrations was accomplished with deionized, double-distilled water. The stock free chlorine (HOCl) solution was prepared from 4% sodium hypochlorite (NaOCl) (Aldrich), diluted to approximately 4000 mg L1 as Cl2, and stored in an aluminum foilcovered glass-stoppered flask. The stock solution was periodically standardized by DPD/FAS titration (Standard Methods for the Examination of Waste and Wastewater, 1998). The glycine solution, used as a precursor to yield ClCN from chlorination, was prepared from a stock solution (100 mg L1 as N) obtained from a reagent-grade chemical (Sigma). Sodium thiosulfate (Na2S2O3) (Fluka), sodium sulfite (Na2SO3) (Nacalai Tesque), sodium metabisulfite (Na2S2O5) (Sigma), and ferrous chloride (FeCl2) (Riedel-deHae¨n) were freshly dissolved into dilution water to make stock solutions of predetermined concentrations. Precisely weighted zero-valent iron (Fe0) powders (Connelly-GPM ETICC-1004) were directly applied to a test solution at predetermined solid-to-liquid ratios (concentrations). Due to the prohibition on importing ClCN to Hong Kong and to the short shelf-life of ClCN, standards of ClCN for calibration curve establishment were freshly, in situ developed by 30-min chlorination of glycine solution (0.5 mg L1 as N) at a free chlorine dosage of 5 mg L1 (as Cl2) and phosphate-buffered pH of 7. This protocol had demonstrated to yield consistent aqueous ClCN concentrations of 900 mg L1 while leaving no residual chlorine in the solution (Shang et al., 2000). Thereafter, dilutions were made to give standards at 900, 100, 10, and 2 mg L1. 2.2. Analytical methods The membrane introduction mass spectrometry (MIMS) method was applied for on-line, rapid aqueous ClCN monitoring and analysis. MIMS employs a membrane interface to introduce an aqueous solution, without any sample work-up, directly to a mass spectrometer by allowing the ClCN to permeate continuously through the membrane while water and ionic species are rejected (Shang et al., 2000; Pederson et al., 1999; Yang and Shang, 2004, 2005). The detailed description of the MIMS method for ClCN quantification can be found in Yang and Shang (2005). In brief, it
ARTICLE IN PRESS 2116
C. Shang et al. / Water Research 39 (2005) 2114–2124
was based on a modification of an HP 5892 bench-top GC/MS (Hewlett-Packard) containing an HP 5972A mass selective detector (MSD) using electron ionization (70 eV) and the selected ion monitoring (SIM) mode. ClCN concentrations were directly determined by comparison of the ion abundance measurements at m/z 61 amu (35ClCN+) with those developed from a series of standard solutions, as the isotopic ion (37ClCN+) at m/z 63 amu may interfere with the molecular ion of an intermediate, N-chloromethanimine (CH2N35Cl+) (Yang and Shang, 2005). The limit of detection of ClCN was 1 mg L1. 2.3. Experimental procedures Most experimental runs were conducted in a 0.5-L well-mixed (using a magnetic stirrer and a stirring bar), parafilm-sealed glass reactor with a 500 mL aqueous glycine solution (0.5 mg L1 as N), phosphate-buffered to pH 5, 6.9, or 9 (0.01 M) at room temperature. The solution pH was periodically checked and no change of the pH was observed in all cases. The glycine solution was first pumped to the MIMS system for 5 min, at which time NaOCl (5 mg L1 as Cl2) was added (the first arrow in all figures) for 30 min to synthesize ClCN. The concentration of ClCN increased with time until it reached a stable, maximum ClCN yield approximately 20 min after the chlorine addition. Thereafter, the concentration of ClCN remained relatively constant until a specific dechlorination reagent was introduced to the testing solution at 35 min (the second arrow in all figures). At the time of introducing the dechlorination reagent, no residual chlorine remained in the solutions. The introduction of the dechlorination reagent required that the parafilm seal be broken and the reactor be immediately resealed. The ClCN concentration was continuously monitored by MIMS until the test run terminated at 70 min. A few test runs were repeated in a 1-L well-mixed, head-space-free, Teflon bag reactor to confirm that the small head space in the glass reactor and breakage of the parafilm seal unaffected the test results. In the cases of using Fe0 powders, a membrane filter with a pore size of 0.45 mm was installed in the sampling line to prevent damage or fouling of the membrane interface from the insoluble iron.
3. Results and discussion ClCN destruction and its kinetics by the three sulfur compounds at pH 5, 6.9, and 9 were first evaluated. Dechlorination processes using reducing sulfur compounds have been well developed to eliminate chlorine residuals and can be easily applied with commercially available packages. Concentrations of the reducing sulfur compounds used in this study were much higher
than those in the practical range to establish the kinetic rate constants. ClCN destruction with the sulfur compounds was assumed to follow the common dechlorination principle that the chlorine atom of ClCN is dechlorinated and reduced to chloride ions (Cl), as S2 O2 3 (or other sulfur compounds) is oxidized to SO2 : 4 k
2 2 2 S2 O2 3 ðSO3 or S2 O5 Þ þ ClCN ! SO4 þ Cl þ other products;
ð1Þ
where k is the second-order reaction rate constant, by assuming a second-order bimolecular reaction between ClCN and the sulfur compound employed in Eq. (1). The rate of ClCN destruction can be then expressed as d½ClCN ¼ k½ClCN ½reducing sulfur compound dt
(2)
where [i] denotes the molar concentration of compound i and t is time. The destruction of ClCN through the reaction was further characterized by assuming pseudofirst-order decay at high concentrations of the sulfur compounds. As such, the rate of ClCN destruction can be expressed as d½ClCN ¼ kobs ½ClCN ; dt
(3)
where kobs is the observed pseudo-first-order rate constant and is defined as the product of the secondorder rate constant (k) and the initial concentration of the reducing sulfur compound. The kobs therefore can be obtained by the slope of the Ln[ClCN] versus t plot. ClCN is also known to hydrolyze in alkali solution according to the general base-catalyzed reaction (Pedersen and Marin˜as, 2001): kh
ClCN þ H2 O ! HOCN þ Hþ þ Cl
(4)
in which kh is the overall hydrolysis rate constant defined as B kh ¼ koh þ kOH h ½OH þ kh ½B ,
koh ,
kOH h ,
(5)
kBh
and are the water, hydroxide ion, where and generic base-B-assisted rate constants. The koh ranging from 3.29 107 to 2.58 106 s1 and the kOH ranging from 3.5 to 13 M1 s1 have been reported h in the literature (Pedersen and Marin˜as, 2001). Thus, the kobs can be rewritten as kobs ¼ kred þ kh ,
(6)
where kred is the pseudo-first-order rate constant contributed only from the chemical reduction reaction. Figs. 1–3 present the results of the experimental runs using sodium thiosulfate, sodium sulfite, and sodium metabisulfite, respectively. Table 1 lists all calculated
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124 900
pH 5.0
2117
Ion 61 (CN35Cl+•)
800
Conc. of CNCl (µg/L)
700 600 500
1.14 mM Na2S2O3
400 300 200 100
11.4 mM Na2S2O3 0 10.00
20.00
30.00
(a)
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500
1.14 mM Na2S2O3
400 300 200 100
11.4 mM Na2S2O3
0 10.00
20.00
30.00
(b)
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 40
pH 9.0
Conc. of CNCl (µg/L)
35 30 25 20 15
1.14 mM Na2S2O3
10 5
11.4 mM Na2S2O3
0 10.00
20.00
30.00
(c)
40.00
50.00
60.00
70.00
Time (minute)
Fig. 1. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium thiosulfate. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O3.
rate constants and R2 corresponding to the experimental results obtained with variations in species and concentrations of the reducing sulfur compounds and varying
pH levels. As listed in Table 1, in all cases, an R2 of 0.946 or higher is obtained, indicating that the pseudo-firstorder assumption is valid.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2118 900
Ion 61 (CN35Cl+•)
pH 5.0
800
Conc. of CNCl (µg/L)
700 600 500
0.633 mM Na2SO3
400 300 200
1.266 mM Na2SO3
100 0 10.00
20.00
(a)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500 400 300 200
0.633 mM Na2SO3 1.266 mM Na2SO3
100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
70.00
Time (minute)
40
Ion 61 (CN35Cl+•)
pH 9.0
Conc. of CNCl (µg/L)
35 30 25 20 15 10
0.378 mM Na2SO3
5
1.266 mM Na2SO3 0 10.00
(c)
20.00
30.00
40.00
50.00
60.00
70.00
Time (minute)
Fig. 2. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium sulfite. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2SO3.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124 900
pH 5.0
2119
Ion 61 (CN35Cl+•)
800
Conc. of CNCl (µg/L)
700 600 500
0.316 mM Na2S2O5
400
1.266 mM Na2S2O5
300 200 100 0 10.00
20.00
(a)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 900
pH 6.9
800
Conc. of CNCl (µg/L)
700 600 500 400 300 200
0.316 mM Na2S2O5 1.266 mM Na2S2O5
100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
70.00
Time (minute)
Ion 61 (CN35Cl+•) 40
pH 9.0 0 mM Na2S2O5
Conc. of CNCl (µg/L)
35 30 25 20 15 10 5
1.266 mM Na2S2O5
0.316 mM Na2S2O5
0 10.00
(c)
20.00
30.00
40.00
50.00
60.00
70.00
Time (minute)
Fig. 3. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating with sodium metabisulfite. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O5.
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2120
Table 1 Pseudo-first-order rate constant (kobs), second-order rate constant (k), and R2 for regression of ClCN destruction with the three reducing sulfur compounds at pH 5, 6.9, and 9
(a) Na2S2O3
pH
Initial conc. (mM)
Calculated Kobs (min1)
R2
5.0
1.14 11.4 1.14 11.4 1.14 11.4
0.04 0.42 0.04 0.41 0.06 0.42
0.998 0.993 0.999 0.986 0.999 0.991
0.61 0.61 0.61 0.61 0.89 0.61
0.633 1.266 0.633 1.266 0.380 1.266
0.03 0.06 0.14 0.26 0.30 0.87
0.999 0.999 0.999 0.991 0.989 0.986
0.81 0.79 3.66 3.36 13.1 11.4
0.316 1.266 0.316 1.266 0.316 1.266
0.01 0.02 0.11 0.34 0.52 1.95
0.994 0.967 0.994 1.000 0.957 0.946
0.62 0.31 5.60 5.26 27.2 25.7
6.9 9.0 (b) Na2SO3
5.0 6.9 9.0
(c) Na2S2O5
5.0 6.9 9.0
3.1. Dechlorination with sodium thiosulfate Fig. 1(a–c) present the results of ClCN destruction by sodium thiosulfate at pH 5, 6.9, and 9, respectively. The thiosulfate additions varied at 1.14 and 11.4 mM. Generally, although there was a reaction, the carbonbonded chlorine reacted slowly with thiosulfate. Only thiosulfate addition at a high, unpractical concentration (11.4 mM) completely destroyed ClCN in 15 min at pH 6.9. Similar trends were found when the experiments were conducted at buffered pH of 5 and 9 while other conditions remained unchanged. It should be noted that, at pH 9, a sudden drop in the concentration of ClCN was observed right after the addition of the thiosulfate (as shown in Fig. 1(c)). The sudden drop was initially postulated to be due to the breakage of the seal for injecting the thiosulfate. However, the hypothesis was denied by the later confirmation tests using a head-space free, sealed Teflon bag reactor and syringe-injection of the sulfur compound where the sudden drop remained. However, we are unable to explain the observation and such a sharp decrease of ClCN was excluded from the data analysis, considering that the later stage is the key stage for ClCN destruction. Table 1(a) presents the calculated rate constants of ClCN destruction by thiosulfate. At pH 5 and 6.9, the calculated second-order rate constants were identical (0.61 M1 s1) regardless of the variations of thiosulfate concentrations and pH. The result indicates that the same mechanism (ClCN destruction mainly by thiosulfate) is applied and ClCN hydrolysis can be neglected at
Calculated k (M1 s1)
the pH range. In other words, kobsEkred. Using the 6 1 largest values of the koh and the kOH s and h , 2.58 10 1 1 13 M s , respectively, reported in the literature (Pedersen and Marin˜as, 2001), the calculated kh is at least two orders of magnitude smaller than the kred at pH 7 or less. On the other hand, considerable discrepancy was observed at pH 9 wherein k’s of 0.89 and 0.61 M1 s1 were obtained at initial thiosulfate concentrations of 1.14 and 11.4 mM, respectively. The discrepancy can be explained by the extent of the incorporation of ClCN hydrolysis in the destruction process. The relative significance of ClCN hydrolysis on ClCN destruction depends on the relative time scale of the reactions. At high initial thiosulfate concentration (11.4 mM), the dechlorination reaction was fast and ClCN completely disappeared in less than 10 min so that the reaction rate constant was not affected by the ClCN hydrolysis. However, ClCN hydrolysis greatly enhanced the ClCN destruction at low thiosulfate addition (1.14 mM) and pH 9. The average kh at pH 9 was calculated from subtraction (based on Eq. (6)) to be 0.01 min1. The somewhat higher kh at pH 9 obtained in the current study is due to the use of a phosphate buffer. Phosphate at moderate pH (X6.5) has been reported to give an increase in the rate of ClCN hydrolysis (Edwards et al., 1986). 3.2. Dechlorination with sodium sulfite Similar experimental runs were conducted using sodium sulfite (0.633 and 1.266 mM) (see Fig. 2). After
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
dissolution, Na2SO3 is expected to yield an equilibrium 2 mixture of bisulfite (HSO 3 ) and sulfite (SO3 ) in aqueous solution according to its acidity constant (Benjamin, 2001): Ka
2 þ HSO 3 2 H þ SO3 ;
pK a ¼ 7:3.
(7)
As such, the reduction of ClCN is attributable to the dechlorination reactions by bisulfite, sulfite or both based on the solution pH. ClCN hydrolysis may also contribute to the overall ClCN reduction. As shown in Fig. 2, although trends that ClCN destruction rates increased with increases in sodium sulfite dosage and pH were observed, the magnitude of the rate increases with increasing pH was much greater than those obtained with sodium thiosulfate additions and the increase also occurred even at pH changes from 5 to 6.9, at which the ClCN hydrolysis could be neglected. The rate constant calculation (see Table 1(b)) followed the same manner as described above and the second-order rate constants at specific pH agreed reasonably well when comparing the values obtained at different initial dosages. The significant rate increases (from approximately 0.8 to 12 M1 s1) with increasing pH were further demonstrated in the rate constant calculation, indicating that ClCN destruction by sodium sulfite is much more sensitive to the solution pH. To shed some light on the causes, calculation on the overall rate constant expressed as the sum of the rates of reactions with sulfite and bisulfite (modified from Croue and Reckhow, 1989) was adopted: 2 kobs ¼ k1 ½HSO 3 þ k2 ½SO3 þ kh ,
(8)
where the kh can be neglected since kh5kobs and k1 and k2 can be obtained as 0.80 and 12.3 M1 s1 at pH 5 and 9, respectively, when the contribution of one or the other is negligible due to its extremely low concentration. The substantial difference in k1 and k2 is expected since sulfite is a much stronger nucleophile compared to bisulfite (Croue and Reckhow, 1989). Correspondingly, the calculated overall second-order rate constant (k) at pH 6.9 is 4.06 M1 s1, which is not far from the experimental data shown in Table 1(b). These findings suggest that the speciation of S(IV) compounds, rather than ClCN hydrolysis, contributes to the variation in the destruction rates. Compared with ClCN destruction by thiosulfate (see Fig. 1 and Table 1), at pH 6.9 and 9, the dose requirement of sodium sulfite to achieve the same degree of ClCN reduction was much smaller and the secondorder rate constants of ClCN destruction by sodium sulfite were much higher than those of thiosulfate. However, the difference was not noticeable at pH 5, where mainly bisulfite was predominant, indicating that bisulfite is as weak as thiosulfate with regard to ClCN dechlorination. These findings only partially agree with
2121
the literature (White, 1999) that S(IV) compounds are better than thiosulfate for dechlorination. It suggests that only S(IV) compounds in the form of sulfite (SO2 3 ) are effective for destruction of ClCN (presumably other DBPs as well) and controlling the pH at neutral or alkali conditions to yield a higher fraction of sulfite is essential. In addition, providing sufficient contact time is also required. To evaluate the effect of residual sulfite/bisulfite on dissolved oxygen (DO) levels, DO was measured at a sodium sulfite dosage of 1.266 mM and pH 7; and DO reduction from 9 to 3.4 mg L1 was observed at the end of the test run. 3.3. Dechlorination with sodium metabisulfite Sodium metabisulfite (0.316 and 1.266 mM) was tested under similar experimental conditions. As shown in Fig. 3 and Table 1(c), similar to the results of using sodium sulfite, large enhancements in ClCN reduction were observed with increases in metabisulfite additions and pH. Again, the second-order rate constants of the S(IV) compound at specific pH agree reasonably well at different initial dosages. Nevertheless, at the same dosage, ClCN reduction by sodium metabisulfite was notably faster (close to double) than that achieved by sodium sulfite at pH 6.9 and 9. This finding can be explained from the equilibrium between bisulfite and metabisulfite (Yiin et al., 1987): keq
2 2HSO 3 2 S2 O5 þ H2 O;
keq ¼ 6:5 102 M1
at 25 C:
ð9Þ
At the metabisulfite concentrations used (0.316 and 1.266 mM), the equilibrium should produce mainly sulfite/bisulfite in stoichiometric quantities with negligible concentrations of metabisulfite and there should be an equilibrium (Eq. (7)) between sulfite and bisulfite according to the solution pH. To directly obtain the hydrolysis rate constant, kh, at pH 9, a side experiment was conducted without adding any reducing sulfur compound. As shown in Fig. 3, after 35 min, approximately 25% of the produced ClCN was lost from the hydrolysis. By assuming a first-order reaction, the kh was estimated to be 0.01 min1, which was in agreement with that previously estimated from subtraction. Interference at the second highest peak of ClCN (corresponding to CN37Cl+ at m/z 63 amu) concurrent with the presence of a considerable abundance of ions at m/z 65 amu was also observed in all cases. The abundance at m/z 65 amu was attributed to the formation of an intermediate, N-chloromethanimine (CH2N37Cl) (Yang and Shang, 2005). As shown in Fig. 4 as an example of using Na2S2O5, a relatively much sharper decrease in the abundance of m/z 65 amu
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2122
Abundance
Ion 61.00 (60.70 to 61.70): 200B.D CN35Cl+•
5500 5000 4500 4000 3500 3000 2500 2000 1500 1000 500 10.00
20.00
(a)
30.00 40.00 Time-->
50.00
60.00
70.00
Abundance
Ion 65.00 (64.70 to 65.70): 200B.D 350 300 250 200 150 100
CH2N37Cl+•
10.00
20.00
30.00
(b)
40.00
50.00
60.00
70.00
Time-->
Fig. 4. On-line selected ion monitoring of chlorinating/dechlorinating glycine solution (0.5 mg L1 as N) at pH 6.9. Ion chromatogram presented are (a) m/z 61 amu (CN35Cl+) for ClCN and (b) m/z 65 amu (CH2N37Cl+) for CH2NCl. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Na2S2O5 (1.266 mM).
was observed immediately after reducing sulfur compound additions. This finding indicates that this nitrogen-bonded chlorine reacts with reducing sulfur compounds much more quickly than the carbon-bonded chlorine in ClCN. Therefore, the nitrogen–chlorine bond is more vulnerable to breakage than the carbon– chlorine bond during the dechlorination process. Bond dissociation energies of 333.979.6 and 397 729 kJ mol1 for the nitrogen–chlorine bond and the carbon–chlorine bond, respectively, at 25 1C are reported by Kerr (2000). In addition, the half-life of N-chloromethanimine (approximately 1 min) obtained from this study during dechlorination was consistent with that of the organic chloramines (on the order of sub-minutes to a few minutes) reported in the literature (Helz and Nweke, 1995; Jensen and Helz, 1998; Jameel and Helz, 1999) at similar S(IV) dosages. Therefore, when carbon-bonded chlorine is required to be effectively destructed by a common dechlorination process using the reducing sulfur compounds, the condition should be likely able to dechlorinate organic chloramines. Further investigation is needed to verify this assertion.
Shang, 2005; Gordon et al., 2002; Orth and Gillham, 1996). Fe2+ and Fe0 were postulated to be capable of reducing ClCN and experimental trials have been conducted. The trials were managed in the same manner as the previous cases except Fe2+ at 50 and 1000 mg L1 and Fe0 at 10, 50, and 100 g L1 were used at pH 6.9. Figs. 5a and b present the resulting ClCN concentrations treated by Fe2+ and Fe0, respectively. As shown, after 35 min, no ClCN could be reduced by up to 1000 mg L1 of Fe2+, while only approximately 50% of ClCN could be reduced by 100 g L1 Fe0. For comparison, 15 mg L1Fe2+ and 50 g L1 of the same Fe0 powders were found to reduce bromate from 50 to 10 mg L1 within 15 min (Siddiqui et al., 1994) and from 100 to 10 mg L1 within 10 min (Xie and Shang, submitted), respectively. ClCN destruction was found to be more feasible from using S(IV) compounds than Fe2+ in this study, which was contradictory to the opposite tendency of bromate destruction from S(IV) compounds and Fe2+ documented in the literature (Gordon et al., 2002). No conclusive explanation for this can be offered here.
3.4. Dechlorination with ferrous ions and zero-valent iron
4. Conclusions
Both ferrous ions (Fe2+) and zero-valent iron (Fe0) are powerful reducing agents and can effectively remove bromate (BrO 3 ) and chlorinated compounds (Xie and
Reducing sulfur compounds may be effective in eliminating ClCN in aqueous solution. In all cases tested, the ClCN destruction was primarily attributable
ARTICLE IN PRESS C. Shang et al. / Water Research 39 (2005) 2114–2124
2123
Ion 61 (CN35Cl+•) 900
1000 mg/L Fe2+
Fe2+ 50 mg/L Fe2+
Conc. of CNCl (µg/L)
800 700 600 500 400 300 200 100 0 10.00
20.00
30.00
40.00
50.00
60.00
Time (minute)
(a)
Ion 61 (CN35Cl+•) 900
Fe0
10 g/L Fe0
Conc. of CNCl (µg/L)
800
50 g/L Fe0
700 600
100 g/L Fe0
500 400 300 200 100 0 10.00
20.00
(b)
30.00
40.00
50.00
60.00
Time (minute)
Fig. 5. ClCN concentrations of 30-min chlorinating aqueous glycine solutions (0.5 mg L1 as N) followed by 35-min dechlorinating by ferrous ions or zero-valent iron at pH 6.9. First arrow: addition of NaOCl (5 mg L1 as Cl2); second arrow: addition of Fe2+ (or Fe0).
to the chemical reduction pathway and followed secondorder reaction kinetics, though the second-order rate constants varied substantially from approximately 0.3 to 25.7 M1 s1 under different experimental conditions. Hydroxide-assisted ClCN hydrolysis was only significant at pH 9 and also when the observed chemical reduction rate was relatively slow. Among the tested compounds and pH, the S(IV) compound in the form of sulfite (SO2 3 ) displays the best destruction rate. However, the observed slower ClCN reduction rates, compared with those of organic chloramines, suggest that ClCN is likely to contribute to effluent toxicity in current wastewater chlorination/dechlorination practices. Therefore, applying a moderately higher amount of the S(IV) compounds (either sulfite or metabisulfite) at neutral or alkali pH and/or an effluent holding tank to increase the contact time before final discharge may be the possible solution(s) to achieve complete ClCN destruction. Nevertheless, it should be noted that the dechlorination of free chlorine and chloramines has been
found to increase with decreasing pH (Yiin et al., 1987; Yiin and Margerum, 1988). An elevated pH may slow down the removal of free chlorine and inorganic and organic chloramines, which are expected to be the major fractions of chlorine after wastewater chlorination. Balance should be carefully sought. On the other hand, slow ClCN destruction is expected even with a small dosage of the three sulfur compounds, indicating that the common practice of using reducing sulfur compounds (e.g., Na2S2O3) to quench chlorination reactions before analysis may cause significant ClCN losses (presumably losses of other DBPs as well), if long-term preservation and storage are not properly avoided.
Acknowledgements The authors are grateful to the Research Grants Council, Hong Kong for financial support of the project
ARTICLE IN PRESS 2124
C. Shang et al. / Water Research 39 (2005) 2114–2124
under project no. DAG00/01EG01 and to Professor Blatchley III ER of the Purdue University, School of Civil Engineering for his input in developing the MIMS method and at the early stage of the preliminary trials.
References American Conference of Governmental Industrial Hygienists, 1991. Documentation of the Threshold Limit Values ad Biological Exposure Indices, sixth ed. Cincinnati, OH. Benjamin, M.M., 2001. Water Chemistry. Mc-Graw Hill, New York. Croue, J.P., Reckhow, D.A., 1989. Destruction of chlorination byproducts with sulfite. Environ. Sci. Technol. 23 (11), 1412–1419. Edwards, J.O., Erstfeld, T.E., Ibne-rasa, K.M., Levey, G., Moyer, M., 1986. Reaction rates for nucleophiles with cyanogen chloride: comparison with two other diagonal carbon compounds. Int. J. Chem. Kinetics 18, 165–180. Gordon, G., Gauw, R.D., Emmert, G.L., Walters, B.D., Bubnis, B., 2002. Chemical reduction methods for bromate ion removal. J. Am. Water Works Assoc. 94 (2), 91–98. Helz, G.R., Nweke, A.C., 1995. Incompleteness of wastewater dechlorination. Environ. Sci. Technol. 29 (4), 1018–1022. Helz, G.R., Uhler, A.D., Sugam, R., 1985. Dechlorination and trihalomethane yields. Bull. Environ. Contam. Toxicol. 34, 497–503. Hirose, Y., Maeda, N., Ohya, T., Nojima, K., Kanno, S., 1988. Formation of cyanogen chloride by the reaction of amino acids with hypochlorous acid in the presence of ammonium ion. Chemosphere 17 (5), 865–873. Jameel, R.H., Helz, G.R., 1999. Organic chloramines in disinfected wastewaters: Rates of reduction by sulfite and toxicity. Environ. Toxicol. Chem. 18 (9), 1899–1904. Jensen, J.S., Helz, G.R., 1998. Rates of reduction of Nchlorinated peptides by sulfite: Relevance to incomplete dechlorination of wastewaters. Environ. Sci. Technol. 32 (4), 516–522. Kerr, J.A., 2000. Strengths of chemical bonds. In: Lide, D.R. (Ed.), CRC Handbook of Chemistry and Physics 1999–2000, 81st ed. CRC Press, Boca Raton, Florida. Kononen, D.W., 1988. Acute toxicity of cyanogen chloride to Daphnia magna. Bull. Environ. Contam. Toxicol. 41 (3), 371–377. Krasner, S.W., McGuire, M.J., Jacangelo, J.G., Patania, N.L., Reagan, K.M., Aieta, E.M., 1989. The occurrence of disinfection by-products in US drinking water. J. Am. Water Works Assoc. 81 (8), 41–53. Lewis Sr., R.J., 1996. Sax’s Dangerous Properties of Industrial Materials, ninth ed. Van Nostrand Reinhold, New York, NY. Morlay, C., de Laat, J., Dore´, M., Courtois, Y., Houel, N., Montiel, A., 1991. Effect of an addition of sodium-sulfite on the mutagenicity of chlorinated solutions of aquatic humic substances. Bull. Environ. Contam. Toxicol. 47, 15–22.
Ohya, T., Kanno, S., 1985. Formation of cyanide ion or cyanogen chloride through the cleavage of aromatic rings by nitrous acid or chlorine. VIII. On the reaction of humic acid with hypochlorous acid in the presence of ammonium ion. Chemosphere 14, 1717–1722. Ohya, T., Kanno, S., 1989. Formation of cyanide ion or cyanogen chloride through the cleavage of aromatic rings by nitrous acid or chlorine. XI. On the reaction of purine bases with hypochlorous acid in the presence of ammonium ion. Chemosphere 19 (12), 1835–1842. Orth, W.S., Gillham, R.W., 1996. Dechlorination of trichloroethene in aqueous solution using Fe-0. Environ. Sci. Technol. 30 (1), 66–71. Pedersen III, E.J., Marin˜as, B.J., 2001. The hydroxide-assisted hydrolysis of cyanogen chloride in aqueous solution. Water Res. 35 (3), 643–648. Pederson III, E.J., Urbansky, E.T., Marin˜as, B.J., Margerum, D.W., 1999. Formation of cyanogen chloride from the reaction of monochloramine with formaldehyde. Environ. Sci. Technol. 33 (23), 4239–4249. Robinson, J.P., 1967. Chemical warfare. J. Sci. 3, 33–40. Shang, C., Gong, W-L., Blatchley III, E.R., 2000. Breakpoint chemistry and volatile byproduct formation resulting from chlorination of model organic-N compounds. Environ. Sci. Technol. 34 (9), 1721–1728. Siddiqui, M., Amy, G., Ozekin, K., Zhai, W., Westerhoff, P., 1994. Alternative strategies for removing bromate. J. Am. Water Works Assoc. 86 (10), 81–96. Standard Methods for the Examination of Water and Wastewater, 1998. Standard Methods for the Examination of Water and Wastewater, 20th ed. APHA, AWWA, and WEF, Washington, DC. US Environmental Protection Agency (USEPA), 1996. Fed. Regist. 61, 24353–24388. White, G.C., 1999. Handbook of Chlorination and Alternative Disinfectants, fourth ed. Wiley, New York. Woodard, F., 2001. Industrial Waste Treatment Handbook. Butterworth-Heinemann, Woburn, MA. Xie, L., Shang, C. Bromate reduction by zero valent iron in batch tests. J. Am. Water Works Assoc., submitted for publication. Yang, X., Shang, C., 2004. Chlorination byproduct formation in the presence of humic acid, model nitrogenous organic compounds, ammonia, and bromide. Environ. Sci. Technol. 38 (19), 4995–5001. Yang, X., Shang, C., 2005. Quantification of aqueous cyanogen chloride and cyanogen bromide in environmental samples by MIMS. Water Res., in press. Yiin, B.S., Margerum, D.W., 1988. Kinetics of hydrolysis of the chlorosulfate ion. Inorg. Chem. 27, 1670–1672. Yiin, B.S., Walker, D.M., Margerum, D.W., 1987. Non-metal redox kinetics: general-acid-assisted reactions of chloramine with sulfite and hydrogen sulfite. Inorg. Chem. 26, 3435–3441. Zhang, A., Dzombak, D.A., Luthy, R.G., 2004. Formation of free cyanide and cyanogen chloride from chloramination of publicly owned treatment works secondary effluent: laboratory study with model compounds. Water Environ. Res. 76 (2), 113–120.