!! r
~
A
B
n.crd
2
Ix,
O b (
r,
e
12
4
0
12
Time (min) Figure 6. Recorder tracings of Se analysis of a river water sample: (A) Se(-ll,O) plus Se(lV) (procedure 2); (B) total Se (procedure 3); (a) 4,6-dibromopiazselenol tilled water and artificial seawater from the peak height of the piazselenol. Analysis of Simulated Mixtures of Se(O), Se(IV), and Se(V1). Distilled water and artificial seawater containing 4.6 ng of Se(O), 4.8 ng of Se(IV), and 5.9 ng of Se(V1) were analyzed by procedures l, 2, and 3, and good recoveries were found (Table I). Artificial seawater was prepared by dissolving 14.0 g of NaCl, 6.3 g of MgC12.6H20,0.9 g of CaC12.2H20, and 0.4 g of KCl in 500 mL of distilled water (Kall's artificial seawater). Na2S04 was omitted, since it was contaminated by selenite. Accuracy and Precision. The accuracy of the procedure was evaluated by recovery experiments, in which known amounts of Se(0) and Se(V1) were added to seawater, and the samples were analyzed by procedures 2 and 3. In all cases, good recoveries were obtained within experimental error (Table 11). The precision of procedures 2 and 3 was evaluated by analyzing eight samples of seawater (500 and 300 mL, respec-
tively). The mean results obtained from these procedures were 10.4 ng of Se L-l and 51.9 ng of Se L-I with standard deviations of 1.0 and 2.4 and relative standard deviations of 9.6 and 4.6%, respectively. Sample Storage. The total selenium and Se(1V) in seawaters stored in glass bottles were determined every day after sampling. One series of the samples was filtered by a membrane filter (pore size 0.45 pm) and the other was not (Figure 4). Concentrated hydrochloric acid was added to seawaters immediately after sampling (1mL L-l), and one was filtered and the other was not (Figure 5). From these results, it was determined that when seawaters are treated with concentrated hydrochloric acid after sampling (1mL L-l) and filtered, the amount of selenium was not altered within 4 days. Determination of Se in River Water and Seawater. As an example, the gas chromatograms of Se in river water by procedures 2 and 3 are shown in A and B of Figure 6, respectively. Selenium (VI) is calculated from the difference between the peak heights of 4,6-dibromopiazselenol in A and B. The amounts and oxidation states of selenium in river water and seawater in Japan are shown in Tables I11 and IV, respectively. The values are means of three determinations. The amounts of Se(1V) and the total selenium are coincident with values reported previously (5). The amount of Se(-I1,O) in river water is larger than that in seawater, whereas Se(1V) in seawater is more prevalent than in river water. Selenium (VI) is most prevalent in seawater. The total amount of selenium in seawater was 50-70 ng L-I, but in river water the amount varied largely, e.g., from 16 to 230 ng L-I. Literature Cited (1) Chau, Y. K., Riley, J. P., Anal. Chim. Acta, 33,36 (1965). (2) S i l l h , L. G., Soensk Kem. Tidskr., 75,161 (1963). ( 3 ) Sugimura, Y., Suzuki, Y., J . Oceanogr. SOC.Jpn., 33,23 (1977). (4) Yoshii, O., Hiraki, K., Nishikawa, Y., Shigematsu, T., Bunseki Kagaku, 26,91 (1977). (5) Shimoishi, Y., TBei, K., Anal. Chim. Acta, 100,65 (1978). (6) Shimoishi, Y., Talanta, 17,74 (1970).
Received for review October 2, 1979. Accepted January 25,1980. This work was partially supported by a Grant-in-Aid for Environmental Science from the Ministry of Education, Science, and Culture of the Japanese Government, which the authors gratefully acknowledge. Presented at the American Chemical SocietylChemical Society of Japan Congress, Honolulu, Hawaii, April 1 4 , 1979, Division of Environmental Chemistry.
Kinetics of Monochloramine Decomposition in the Presence of Bromide Timothy W. Trofe', Guy W. Inman, Jr.*, and J. Donald Johnson' Department of Environmental Sciences and Engineering, School of Public Health, University of North Carolina, Chapel Hill, N.C. 27514
The electric power industry has traditionally used chlorine as an effective biocide in controlling the growth of microbial slimes and biofouling organisms throughout the cooling water systems of power plants employing either fresh or saline water as a coolant. Many studies have shown that chlorine residuals in discharged cooling waters are toxic to aquatic life forms (1-6) at very low residual chlorine concentrations. The toxicity of these chlorine-induced oxidants depends upon their persistence or stability as well as their chemical form. Present address, Office of Water Research and Technology, Wrightsville Beach Test Facility, Wrightsville Beach, N.C. 28480. Present address, Burroughs Wellcome Co., Greenville, N.C. 27834. 544
Environmental Science & Technology
In order to effectively assess, minimize, and control the impact of chlorine residuals, data are needed on the chemical reactions of chlorine, especially in estuarine and marine waters. The major inorganic reaction pathways of chlorine known to occur in estuarine waters, where chlorine itself does not persist, are shown schematically in Figure 1. Johnson and Inman (7) have studied the kinetics of both the bromide and ammonia reaction pathways at pHs and salinities found in marine waters. In saline waters, ammonia and bromide compete for chlorine and, depending on the salinity, either monochloramine or bromamines may predominate. Although both materials are toxic, the bromamines are more desirable since they are less persistent when discharged into the receiving water system (8,9).Dibromamine (NHBr2) is able to
0013-936X/80/0914-0544$01 .OO/O
@ 1980 American Chemical Society
Monochloramine, a major product formed by chlorine added to natural waters, is a problem because of its persistence and toxicity to marine organisms. One major route for its decomposition is through its reaction with bromide. This reac-
tion was found to obey the rate law -d[NHzCl]/dt = k[H+]. [Br-] [NH2C1],where k = 2.8 f 0.3 X lo6 M-2 s-l at 25 "C. The major reaction product is postulated to be a mixed haloamine, NHBrCl, not previously identified.
disproportionate and enters readily into redox reactions, whereas monochloramine (NHzC1) is a relatively stable oxidant. When monochloramine is formed, its stability is important in determining the toxicity of the water when discharged. Helz (10) has suggested, based on thermodynamic arguments, that monochloramine may, like chlorine, oxidize bromide ion to form monobromamine. Our previous kinetic results (11)indicated that Reaction 1would be more likely to occur.
0.2 "C) solution of KBr and phosphate buffer. A 4-mL aliquot was withdrawn and placed into a 1-cm quartz cell and a series of repetitive scans made from 350 to 210 nm using a Cary Model 219 dual-beam spectrophotometer. Absorbance data were taken at three wavelengths, 220,232, and 245 nm. The 232- and 245-nm wavelengths corresponded to the A,, for dibromamine and monochloramine, while the 220-nm wavefor a previously unreported length was found to be the A,, compound, bromochloramine, NHBrC1, discussed below. Reactant and product concentrations were computed from the absorbance data by solving two simultaneous equations and in two unknowns using the molar absorptivities of "$21 NHBr2 in the initial rate studies and those of NH2Cl and NHBrCl in the rate constant determination. The molar absorptivities of NH&l and NHBrz are given elsewhere (12). Chlorine demand free water (CDFW) was prepared by first deionizing and distilling tap water. Twenty liters was dosed with an appropriate amount of Fisher laboratory grade NaOCl (4-6%) to produce a 3 mg/L residual as free chlorine. The water was kept stoppered for 2-3 days and then dechlorinated by ultraviolet radiation to remove the remaining chlorine residual. Chlorine demand free water was used as a diluent in all reagents and reaction solutions. Monochloramine stocks (25.0 mM) were prepared by mixing an appropriate amount of NaOCl stock (its concentration calculated from the absorbance of OC1- at 292 nm) to a solution of NH4Cl buffered a t pH 8.75. The ammohia to chlorine molar ratio was maintained a t 2:l. The monochloramine stock concentration was determined spectrophotometrically from the absorbance a t 245 nm prior to each experiment. Stock potassium bromide solutions were prepared by the addition of Fisher Certified ACS grade potassium bromide in demand free water and dilution to volume. A potassium phosphate buffer stock (0.20 M in Pod3-) was prepared by adding enough 0.2 M K2HP04 to 0.2 M KH2P04 until the desired pH was obtained. In order to remove trace amounts of reducing agents, the stock was filtered (0.45 pm), chlorinated, and irradiated with UV light. The temperature was controlled ( 2 5 f 0.2 "C) by connecting a Lauda-Brinkman Model K-S/RD circulating water bath to a water-jacketed beaker and to the water jacket of the Cary 219 spectrophotometer. Artificial Seawater Kinetic Studies. The reaction was also studied in artificial seawater solutions where the salinity was 10%0and the bromide concentration was 25.0 mM. Artificial seawater stocks were prepared according to Sillen's formula ( 1 3 ) .The initial monochloramine concentration was 0.05 mM (3.5 mg/L as Cl2). The solutions were phosphate buffered, 7.0 mM at pH 6.85 and 2.0 mM at pH 7.25. This was done to prevent magnesium salt precipitation at the higher pH. The runs were carried out using a 5:l molar ratio of Br-/NH&l. Seawater collected near Grand Bahama Island ( S = 36.5%0)was filtered through a 0.45-pm filter and diluted with CDFW to 17%0salinity. Reactant and product concentrations were calculated as described in the initial rates section, where 4-cm quartz cells were used instead of 1-cm cells. To ensure that the reaction rates at these low concentrations were not affected by UV irradiation, the sample and reference cells in the spectrophotometer were shielded against incident radiation between scans. The pH was monitored in the reaction solutien over a 3-h time period. Observed rate constants were obtained from second-order plots of In ([Br-]/[NH&l]) vs. time where k&sd = slope/([Br-I0 - [NH2C1I0).
2NH2C1+ 2Br-
-
NHBrz
+ 2C1-
In this paper we present the results of a study to determine the kinetics and mechanism of monochloramine decomposition in the presence of bromide ion. Both spectral and kinetic evidence show that the expected reaction (1)does not occur, but that a new product, NHBrCl, is produced as monochloramine oxidizes bromide. The rate of this reaction in estuarine waters is important in determining the chemical makeup and stability of chlorine-induced oxidants produced in cooling waters and sewage treatment plant effluents.
Experimental The reaction orders with respect to the concentration of each component were determined by the method of initial rates. Additional experiments were carried out a t lower concentrations, higher ionic strengths, and pHs typical of estuarine waters. Observed rate constants, computed for 90% of complete reaction, were then used to evaluate an overall empirical rate constant. Initial R a t e Studies. The initial concentrations of monochloramine and bromide were varied from 0.050 to 1.2 mM and from 1.00 to 22.5 mM, respectively. The excess ammonia nitrogen, N T ' ~= [NH4+] + [NHs], ranged from 0.010 to 19.0 mM, while the pH was varied between 6.8 and 8.2. These experiments were also conducted under pseudo-first-order conditions where the molar ratio of bromide to monochloramine was greater than 15:l. The initial decomposition rates of monochloramine were determined from the slope8 of the In [NH2Cl] vs. time plots. The reaction was initiated by adding 3 to 4 mL of a 25.0 mM monochloramine stock to a stirred and thermostated (25 f
HOBr
NH ,Br
IHOBr NHBr2
I
HOBr
1 NBr3
Figure 1. Principal inorganic reaction pathways of chlorine in saline waters
Volume 14, Number 5,May 1980 545
4.50
[Bi]' = 16.1mM
4.50
3.50
4.00
3.00
2.50
850
-LOO
Figure 2. Effect of initial monochloramine concentration on initial rate of reaction
3 10
2 IO
2 60
I60
-Log [ B i ]
Flgure 3. Effect of initial bromide ion concentration on initial rate of
reaction
Reaction Product Determination. The solubilities of the three chloramines in ether vary considerably, and this property has been used as a separation method (14). In addition, the bromamines have been formed in ether and their UV spectra characterized (12). The formation of haloamines during the reaction was confirmed by extracting aliquots of the reaction solution into anhydrous ether (J. T. Baker certified ACS grade), removing the ether phase, and scanning the spectrum between 350 and 210 nm. This procedure was also used to obtain reference spectra of monochloramine, dichloramine, monobromamine, and dibromamine ether extracts from pure solutions. Ether extraction was also used to confirm an absorbance peak at 218 nm from bromochloramine (NHBrCl). Results and Discussion Order with Respect to Monochloramine. The reaction order with respect to monochloramine concentration was determined by varying the initial concentration of mono546
7 50
800
- L o p (Nn,Cl]
Environmental Science & Technology
700
650
[H+]
Figure 4. Effect of initial hydrogen ion concentration on initial rate of
reaction chloramine between 0.05 and 1.2 mM NH2C1. The bromide concentration was maintained at 16.0 mM and the solutions were buffered at pH 7.6. Figure 2 shows a plot of the logarithm of the initial rate, v, vs. the logarithm of the initial monochloramine concentration. The slope was evaluated by linear regression as 0.99 f 0.15 for the order with respect to monochloramine. Order with Respect to Bromide Ion. Similarly, the order with respect to bromide ion was investigated by varying the bromide ion concentration from 1.00 to 23.0 mM Br-. The pH was maintained a t 7.6, and the initial concentration of monochloramine was constant at 0.25 mM. Figure 3 is a plot of the logarithm of the initial rate vs. the logarithm of the initial concentration of bromide ion where the slope was evaluated as 1.01 f 0.06. Order with Respect to Hydrogen Ion. The order with respect to hydrogen ion was determined by varying the pH between 6.8 and 8.3 in six kinetic runs. The monochloramine and bromide ion concentrations were held at 0.75 mM NHzCl and 16.0 mM Br-, respectively. Figure 4 is a plot of the logarithm of the initial rate vs. pH. The slope was evaluated as 0.98 f 0.14. The reaction orders obtained from these experiments are consistent with the empirical rate equation:
-d["2C11 = k [NHzCl][Br-] [H+] dt If the hydrogen ion concentration remains constant, Equation 2 can be integrated to its second-order form as follows: In
(a) = k[H+]([Br-]O
- [NH2Cll0)t - In
[NH2C1I0
( [Br-]O )
(3)
Identification of Bromochloramine. During the initial rate experiments and the artificial seawater runs, an absorbance peak near 220 nm was frequently observed in the UV spectrum of the reaction solution. Figure 5 shows the peak appearing 2 h after the reaction was initiated. The absorbance peak at 220 nm could not be explained by the presence of dioccurs at 232 nm (12). Other halobromamine since its A,, amines that absorb in this region are dichloramine and trichloramine, which have absorbance maxima at 206 and 220 nm, respectively (15).The formation of these two compounds
0.08 218 nanometers
0.06 E
e
s
0
4 004
210
250
230
270
Wavelength (nrn)
Figure 5. Series of repetitive scans made on artificial seawater (s%o = 10) dosed with 0.05 mM monochloramineshowing the appearance of an absorbance peak at 220 nm
was not probable, since the total ammonia nitrogen to halogen molar ratio, NT:CIT+,was always greater than one during the course of the reaction. Ether extraction followed by UV absorbance scan of the ether extract was made on the following reaction mixtures in an effort to isolate and characterize the absorbance peak at 220 nm. (a) NH2Cl+ Br(b) HOBr (c) HOC1
+ NH&! + NH2Br
-
Wavelength (nm)
Figure 6. UV absorbance spectrum on the ether extract of a reaction solution of (a) monochloramine and bromide ion, (b) hypobromous acid and monochloramine, (c) hypochlorous acid and monobromamine, and (d) monochloramine and dibromamine. All solution mixtures contained excess ammonia nitrogen
:;;;;etion
L"
peak a t 218 nm
0
NTxg aqueous
I I stoichiometry 2 I stoichiometry
(d) NH2C1+ NHBr2 In each case the ether extract exhibited an absorbance maximum at 218 f 2 nm as shown in Figure 6. The 218-nm absorbance peak was not due to dibromamine, since its , , ,A in ether is 234 nm. Table I was constructed from UV absorbance data for the chloramines and bromamines in aqueous solution. The chloramines exhibit peaks at lower wavelengths than their bromamine counterparts (NHC12 = 206 nm vs. NHBr2 = 232 nm). Considering the compound, NHBrC1, we expect its absorbance maximum to fall somewhere between those of NHC12 and NHBr2. Thus, we postulate that NHBrCl is responsible for the peak near 220 nm in the reaction solution (Figure 5 )* A direct consequence of postulating NHBrCl as the primary reaction product rather than NHBr2 is the effect upon the overall reaction stoichiometry. When dibromamine is the reaction product, Equation 4 gives a 1:l stoichiometry between monochloramine and bromide. Alternatively, when bromochloramine is the reaction product, Equation 5 yields a 2:l stoichiometry. One mole of monochloramine is required to oxidize bromide and an additional mole for substitution to form NHBrCl. Table 1. Maximum Absorbance Wavelengths for the Haloamines ( X = CI, Br) NHPX NHX2 NX3 a
chloramlnes
brornamlnes
245 206, 295 220
278 232,350
258
Some data from ref 12. Wavelengths are reported in nanometers.
20.
c 15.
2 z
I
\ 'L
m
I
J
0 01 40
BO
120
0
Time (minuter)
Figure 7. Second-order plots comparing 1:l stoichiometry for NHBr2 formation to 2: 1 stoichiometry for NHBrCl formation
2NH&1+ 2Br2NH2C1+ Br-
-
-
+ 2C1- + NH3 NHBrCl + C1- + NH3 NHBr2
(4) (5)
To test whether a 1:l or 2:l stoichiometry exists, secondorder conditions were chosen where the concentrations of monochloramine (0.96 mM) and bromide (1.2 mM) were nearly equal. Using the form of the integrated rate equation shown in Equation 3, a plot of In ([Br-]/[NH&l]) vs. time was prepared as in Figure 7 . The logarithmic function (Equation 3) was evaluated using both 1:l and 2:l stoichiometries. The 2:l stoichiometry showed a good fit of the kinetic data throughout the experiment, while the 1:l stoichiometry plot Volume 14, Number 5, May 1980
547
~~~
Table II. Summary of Kinetic Data for Decomposition of Monochloramine in Artificial Seawater a
PH
AWS-M-02 ASW-M-03 ASW-M-04 ASW-M-05 ASW-M-06 BSW-M-03
0.05 0.05
0.25 0.25 0.25 0.25 0.25 0.43
0.05
0.05 0.05 0.05
6.84 7.03 6.72 7.15 7.25 7.85
13.1 10.1 17.7 7.3 6.4 1.56
a Temperature = 25 O C , S = 10%0,I = 0.20. * BSW-M-03 (Bahamas seawater, S = 17'700).
deviated from linearity with time as shown. The bromide concentrations were calculated from both the 1:l and 2:l stoichiometries as shown in Equations 6 and 7 , respectively. The concentration of monochloramine present a t any time during the reaction was calculated from absorbance data taken at 245 and 220 nm. The molar absorptivity of bromochloramine at 220 nm was estimated as 2100 M-l cm-I based upon the known molar absorptivity values for NHC12 and NHBr2 of 2100 and 1900 M-l cm-l. [Br-lt = [Br-IO - [ ~ ~ 2 ~ ~ ] c o n s u m e . d [Br-lt = [Br-I0 - 1/2[~H2C~]consumed
8.50 8.00 7.50 7.00 a
Formation of bromochloramine, NHBrCl, explains both the observed UV absorbance change in the reaction mixture shown in Figure 5 and the stoichiometry of the reaction. Determination of a n Empirical R a t e Constant. Second-order observed rate constants were evaluated using artificial seawater stocks where the salinity was 10%. The molar ratio of bromide to monochloramine was 5:l and the pH varied between 6.8 and 7.2. Temperature was maintained at 25 "C. Table 11lists the observed rate constant, hobs& as a function of pH for the kinetic runs. Run number BSW-M-03 represents an experiment conducted with Bahamian seawater diluted to 17%0salinity with CDFW. Figure 8 is a plot of kobsd vs. hydrogen ion concentration, where the empirical rate constant, h , is given by the slope as 2.8 f 0.3 X lo6 M-2 s-l a t 25 "C. Reaction Mechanism. Experimental evidence has shown
rallnlty, %o 17
35
10
50 15
25
a 2.5 0.75
90 29 9 2.8
5 1.5
5
188 60 20 6
Half-lives reported in hours: temperature = 25 O C .
the reaction to be first order with respect to monochloramine, bromide ion, and hydrogen ion. Also, it has been shown that an overall 2:l stoichiometric relationship exists between monochloramine and bromide. Spectral evidence also shows that a mixed haloamine, bromochloramine (NHBrCl), is a principal reaction product. The following reaction mechanism is consistent with the observed stoichiometry and empirical rate law: NH2Cl+ H+
+ Br-
(6) (7)
~
Table 111. Half-Life Values for a 1.0 mg L-' as CIp Dose of Monochloramine as a Function of pH and Salinity a
NH3Br+
2NH3C1+
ki
+NH3Br+ slow
+ C1-
+ NHzCl+fast NHBrCl + NH4+ R2
(8)
(9)
(10)
Equations 9 and 10 are representative of competitive, consecutive second-order reactions preceded by a rapid equilibrium, Equation 8. This general reaction scheme has been treated in detail elsewhere (16). The following differential equation can be derived directly from Equations 8, 9, and 10:
-d[NH2C11 = 2k1K,[NH2C1][Br-][H+] (11) dt Equation 11 is identical with the empirically derived rate expression, Equation 2, where the empirical rate constant, k , is equal to 2k1KP. Since K, = 28 M-l, according to Gray, Margerum, and Huffman ( 1 7 ) ,kl, the reaction rate constant, equals 5.0 X 104 M-l s-l at 25 "C. This value is approximately ten times that of the corresponding rate constant for the oxidation of bromide by HOCl (8, 18).Although NH3C1+ is a better oxidant than HOCl because OH- is more basic than "3, the observed rate of bromide oxidation by monochloramine is very slow due to the weak basicity of NH2C1.
35L
0
10.0
5.0 [H*]
X
lo',
150
20 0
M
Figure 8. Observed second-order rate constant, kobsd, vs. [H+] for artificial seawater runs and Bahama seawater 548
Environmental Science & Technology
Applications The rate of monochloramine decomposition in estuarine waters is especially important due to the toxicity of monochloramine to marine life (5, 6). In addition to demand and substitution reactions with various inorganic and organic substrates, monochloramine also reacts with naturally occurring bromide ion in saline waters. The rate of monochloramine disappearance by this route is primarily a function of pH and salinity. The reaction is first order with respect to both hydrogen ion and bromide ion. Table I11 lists the half-lives in hours of monochloramine (
