Environ. Sci. Technol. 2004, 38, 2111-2119
Bromamine Decomposition Kinetics in Aqueous Solutions HONGXIA LEI, BENITO J. MARIN ˜ AS,* AND ROGER A. MINEAR Department of Civil and Environmental Engineering, University of Illinois at Urbana-Champaign, 205 North Mathews Avenue, Urbana, Illinois 61801
The objectives of this study are to investigate the kinetics of bromamine decomposition and to identify the corresponding relevant reactions. Experiments were performed with a stopped-flow spectrophotometer system. Experimental variables investigated included pH (6.5-9.5), bromamines concentration (0.15-0.50 mM), ammonia to bromine ratio (5-100), and phosphate and carbonate buffers concentration (5-40 mM). The experimental results were consistent with a reaction scheme that involved the reversible disproportionation of monobromamine into k1
} NHBr2 + NH3), dibromamine and ammonia (2NH2Br {\ k-1 followed by irreversible decomposition of monobromamine k2
and dibromamine into products (2NHBr2 98 products k3
and NH2Br + NHBr2 98 products). The monobromamine disproportionation reaction was found to undergo general acid catalysis, and the two subsequent decomposition reactions were found to experience base catalysis. Experimental results were analyzed for the determination of catalysis terms corresponding to H+, NH4+, H2PO4-, HCO3-, and H2O for rate constants k1 and k-1; HPO42- and H2O for k2; and OH-, CO32-, and H2O for k3. These constants were fitted with the Brønsted relationship, and the resulting fitting expressions were used to calculate any relevant catalysis rate constants that could not be determined at the range of experimental conditions used.
Introduction The application of bromamines for drinking water disinfection was given serious consideration several decades ago mainly because these chemicals were found to be stronger disinfectants than the more stable chloramines (1-3). However, the direct application of bromine-based chemicals for drinking water disinfection became in disuse because subsequent information revealed that brominated disinfection byproducts (DBPs) had relatively high genotoxicity and carcinogenicity. Nevertheless, the presence of bromamines has persisted in drinking water because of their formation from bromide ion during chlorination and ozonation of natural water (4-6). The level of bromamines that could be formed in natural waters could vary widely depending on the concentration of bromide ion (typically ranging from single digit to several hundred micrograms per liter) and other water quality parameters such as ammonia concentration, pH, and natural organic matter content as well as on the applied level of chlorine and ozone. * Corresponding author e-mail:
[email protected]; phone: (217)333-6961; fax: (217)333-6968. 10.1021/es034726h CCC: $27.50 Published on Web 02/21/2004
2004 American Chemical Society
Bromamines affect the formation of DBPs of public health concern such as bromate and cyanogen halides (7, 8). The formation of bromate from reactions initiated by ozone and bromide ion can be reduced in the presence of ammonia or chloramines because hypobromous acid (HOBr), an intermediate species that forms from the reaction between ozone and bromide ion and plays a key role in bromate formation, reacts with ammonia and chloramines forming bromamines, which do not get involved in bromate formation (4). In contrast, chloramines and bromamines can enhance the production of cyanogen chloride and cyanogen bromide through reactions with the common DBP formaldehyde (6, 9, 10). Bromamine chemistry has been shown to have some similarity with the better-understood chloramine chemistry. Inman and Johnson (11) found that mixing HOBr with ammonia (NH3) in the range of pH 7.0-8.4 would first result in the formation of monobromamine (NH2Br) and then NH2Br would disproportionate into dibromamine (NHBr2) according to the reversible reaction: k1
2NH2Br {\ } NHBr2 + NH3 k
(1)
-1
Both NH2Br and NHBr2 underwent subsequent decomposition, but in general NHBr2 appeared to decompose more rapidly. Consequently, the overall rate of bromamine decomposition was slowed by increasing both the initial ammonia to bromine ratio and the pH, with the resulting higher NH3 concentration forcing the equilibrium toward the left side of reaction 1. The forward and reverse rates of reaction 1 were also shown to increase with decreasing pH and increasing phosphate buffer concentration due to the occurrence of acid catalysis (11). In contrast with the observation by Inman and Johnson (11), Cristina et al. (12) reported that the formation of NHBr2 from NH2Br was first-order in NH2Br rather than the second order indicated with reaction 1. No explanation was provided to resolve this apparent discrepancy. Cristina et al. (12) also showed that the overall decomposition of bromamine over the wide pH range of 6-12 followed second-order kinetics when [NH3]T/[HOBr] > 1. However, unexplained deviations from this general trend were observed for the data at the beginning of the experiments performed at relatively high and low pH values within this range. The decomposition of NHBr2 has also been studied by the method of initial rates (13, 14). However, the experimental results were not consistent. The apparent reaction order for the decomposition of NHBr2 was determined to be 2.5 in NHBr2 at pH 6. In contrast, the reaction was found to be second-order in NHBr2 at pH 7 and pH 8. Despite the fact that the chemistry of bromamines has been studied over several decades, it is far from being fully understood. There is no agreement on the chemical kinetics and reaction pathways. The reasons for the discrepancies observed in previous studies need to be elucidated, and a unified bromamine decomposition mechanism should be developed. In view of the research needs identified in the preceding paragraph, the objectives of this study are to investigate the kinetics of bromamine decomposition and identify the corresponding relevant reactions for conditions at which monobromamine and dibromamine are the predominant combined bromine species present in solution. Experiments were performed in organic-free synthetic solutions with ammonia present in stoichiometric excess with respect to VOL. 38, NO. 7, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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bromine. The experimental variables investigated were pH, ammonia to bromine ratio, and phosphate and carbonate buffer concentrations.
Experimental Section Reagents. Distilled deionized (DDI) water was used in all experiments. Hypochlorite solutions were prepared by diluting Fisher reagent-grade sodium hypochlorite (4-6%). The concentrations were determined by spectrophotometry using a molar absorptivity of 362 M-1 cm-1 at λmax of 292 nm for OCl- (15). Bromine solutions with target molar ratios [Br-]/[OBr-] ) 0.05 and [Cl-]/[OBr-] ) 1 were prepared by reacting NaOCl with Br- according to
OCl- + Br- f OBr- + Cl-
(2)
at pH 11, molar Br-/NaOCl ratio of 1.05, and ionic strength of 0.1 M (NaClO4). The reaction was allowed to proceed for 3 d. The final concentration was determined by spectrophotometry using a molar absorptivity of 332 M-1 cm-1 at the λmax of 329 nm for OBr- (16). The desired stock concentration of OBr- was then achieved by diluting the stock solution with a solution with ionic strength of 0.1 M and pH 11.0 prepared by adding NaClO4 and NaOH to DDI water. The resulting OBr- stock solution was kept for up to 3 weeks, a period during which the bromine concentration was found to decrease by less than 4%. Two additional methods for preparing bromine solutions were also followed for the special purpose of assessing potential effects of bromide and chloride ions on bromamine decomposition kinetics. A free bromine solution with target molar ratios [Br-]/[OBr-] ) 1 and [Cl-]/[OBr-] ) 0 was prepared fresh daily by dissolving liquid bromine (∼99%) in a 0.01 N NaOH solution in DDI water and allowing it to react according to
Br2 + H2O f HOBr + Br- + H+
(3)
The pH was then adjusted to 11.0 with a 0.1 or 0.01 N NaOH solution. An additional stock with higher bromide ion concentration [Br-]/[OBr-] ) 4 was also prepared by this method with the additional step of adding the mass of sodium bromide required to achieve this ratio. Free bromine with target molar ratios [Br-]/[OBr-] ) 0 and [Cl-]/[OBr-] ) 0 was prepared by reacting ozone and sodium bromide at respective concentrations of 1 and 0.8 mM according to
O3 + Br- + H2O f O2 + HOBr + OH-
(4)
The conditions used were a pH of 4 and a phosphate buffer concentration of 0.01 M (17). After allowing the reaction to proceed for 24 h, nitrogen gas was bubbled through the solution for 20 min to remove any residual ozone. The pH was then adjusted to 11.0, and the solution was stored at 4 °C until being used within a few days. Bromide or chloride ion solutions were prepared by dissolving NaBr or NaCl in DDI water. Ammonia solutions were prepared by dissolving NH4Cl in DDI water and then adding a 0.2 M stock solution of buffer (phosphate or carbonate) to achieve a target buffer concentration of 0.02 M and NaClO4 to adjust the ionic strength to 0.1 M. The pH was then adjusted to the target value with a NaOH or HCl solution. Ammonia solutions with no chloride ions were prepared by a similar method except using (NH4)2SO4 instead of the chloride salt. Phosphate buffer stock solutions (0.2 M) were prepared by dissolving Na2HPO4 and NaH2PO4 in DDI water, followed by pH adjustment (pH 7). Carbonate buffer solutions were prepared by dissolving NaHCO3 in DDI water 2112
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followed by adjusting the pH with NaOH solutions to target values. All chemicals used were reagent grade. Experimental Methods. A Corning 340 digital pH meter was used for all pH measurements. The measured H+ ion activity {H+} ({H+} ) γH+[H+]) was converted to molar concentration [H+] using an activity coefficient γH+ calculated from the extended Debye-Hu ¨ ckel equation and the Gu ¨ ntelberg approximation (18). In this work, the two values were strictly differentiated with the terms pHc (pHc ) -log[H+]) and pH (pH ) -log{H+}). The dissociation constants (pKa) for all acid-base species used in this research were also corrected for ionic strength (19). A Beckman DU 7500 diode array spectrophotometer (Beckman Instruments, Inc., Fullerton, CA) was used to obtain UV-vis spectra and take measurements at single wavelengths. All kinetic studies were conducted with a SX.18MV stopped-flow reaction analyzer (Applied Photophysics Ltd., Leatherhad, Surrey, U.K.). This instrument allowed fast mixing of equal volumes of two reactants as well as sequential mixing (i.e., mixing two reactants first and then mixing the products with a third reagent). The instrument could be configured to study both relatively fast and slow reactions, with reaction times ranging from milliseconds to 1000 s. The two primary reactants used in all stopped-flow runs were HOBr and NH3. HOBr was prepared at the target concentration by diluting the stock HOBr solution with a solution having an ionic strength of 0.1 M and pH 11.0. NH3 solutions were prepared similarly, by diluting the NH3 stock with a solution having a buffer concentration of 0.02 M (phosphate or carbonate), an ionic strength of 0.1 M, and the same pH of the ammonia stock. Once ammonia was mixed with HOBr, the buffer concentration dropped to 0.01 M, and NH2Br was formed within 1.6 ms at pH 9.5 (20). Bromamine decomposition studies were performed with NH3 in excess, typically in the range of 2-10 mM, and HOBr concentrations in the range of 0.15-0.5 mM. For each set of conditions, the kinetic experiment was run in triplicate at each of the three fixed wavelengths of 278, 232, or 258 nm corresponding to maximum absorbance of NH2Br, NHBr2, or tribromamine (NBr3), respectively. The concentrations of bromamines were then determined by solving the resulting system of three linear expressions for total absorbance using the corresponding molar absorptivities. Ammonia was generally present in excess (Table 1), and thus its concentration did not need to be monitored over time. Experimental Matrix. All experiments performed to study bromamine decomposition kinetics are summarized in Table 1. Ammonia was added in excess concentration with respect to that of total bromine to ensure that monobromamine and dibromamine were the predominant bromine species present in solution. Phosphate or carbonate buffer were used to maintain the solution pH constant. Initial experiments were performed to evaluate the molar absorptivity values for monobromamine (set ABS-1) and dibromamine (set ABS-2) and the effect of bromide ion (set Br-eff). The molar absorptivity values for each combined bromine species could be determined independently under the experimental conditions specified in Table 1 because monobromamine and dibromamine were found to become the predominant combined bromine species present in solution at relatively high and low pH, respectively, within the overall pH range of 6.5-9.5 investigated. The kinetics of bromamine decomposition was investigated under conditions designed to assess catalysis effects due to H+ (sets HN-1-3), ammonium ion (sets NN-1 and NN-2), phosphate ions (sets PP and HP-1-3), and carbonate ions (sets CN-1 and CN-2). Once again, the decomposition of monobromamine and dibromamine was investigated by performing experiments at relatively high and low pH, respectively. An additional set of
TABLE 1. Summary of Experiments Performed To Study Bromamine Decomposition at 25 °C and Ionic Strength of 0.1 M (NaOCl4) exp. set
pH
[HOBr]0 (mM)
CT,NH3,0 (mM)
typeb
buffer concn (M)d
ABS-1 ABS-2 Br-effa HN-1 HN-2 HN-3 NN-1 NN-2 CN-1 CN-2 PP HP-1 HP-2 HP-3 NP
8.81-9.44 6.51, 6.74 6.71, 8.48 9.01 9.24 9.50 8.98 9.44 8.98 9.43 7.21 6.77 7.95 8.81 7.95
0.3-0.5 0.4, 0.5 0.50 0.15-0.50 0.15-0.50 0.15-0.50 0.4 0.4 0.4 0.4 0.5 0.15-0.50 0.15-0.50 0.15-0.50 0.4
5-40 8, 10 10 8.15-8.50 10.01-10.36 13.32-13.67 5-40 5-30 10 10 10 10 10 10 2-20
phosphate/no bufferc phosphate phosphate no bufferc no bufferc no bufferc no bufferc no bufferc carbonate carbonate phosphate phosphate phosphate phosphate phosphate
0.01 0.01 0.01 na na na na na 0.005-0.04 0.005-0.04 0.005-0.04 0.01 0.01 0.01 0.01
a Br- varied from 0 to 2 mM. including ammonia.
b
Buffering also provided by ammonia in all cases. c Ammonia served as default buffer.
experiments (NP) was performed at an intermediate solution pH to compare data to model predictions made with reactions elucidated and rate constants obtained from fitting the previous experiments. Data Analysis. The reaction mechanism for this study was found to involve one reversible reaction and two subsequent decomposition reactions, and at least two of these reactions (i.e., reversible reaction and one of the subsequent decomposition reactions) were always occurring concurrently at a given set of experimental conditions. The corresponding rate constants were calculated by fitting experimental data (monobromamine and dibromamine concentrations as a function of time) sets to the corresponding rate expressions with the program Dynafit (21). This software could be used to fit each data set independently or to fit several data sets simultaneously as described in the Results and Discussion section. In either case, the data was fitted to the rate expressions by nonlinear regression using the LevenbergMarquardt algorithm. A trial-and-error approach was used to assess if all three reactions (i.e., reversible reaction plus both subsequent decomposition pathways) or just two of the reactions (reversible reaction plus one of the two decomposition reactions) were relevant for a given set of experimental conditions.
Results and Discussion NH2Br and NHBr2 Molar Absorptivities. To resolve the overlapping spectra of NH2Br and NHBr2, precise molar absorptivity values for these two bromamines were needed. NH2Br molar absorptivity was determined by mixing known concentrations of HOBr with NH3 with the stopped-flow system for experimental conditions at which NH2Br would be predominant over NHBr2 (set ABS-1). HOBr reacted with NH3 according to
HOBr + NH3 a NH2Br + H2O
(5)
Although there are some discrepancies among values reported for the rate constants of reaction 5, all reported values for the forward reaction are greater than 107 M-1 s-1 (20, 22) and that for the backward reaction is approximately 1.5 × 10-3 M-1 s-1 (22). Accordingly, under the experimental conditions used in this study (set ABS-1, Table 1), practically all of the HOBr was converted to NH2Br in 0.1 ms or approximately one thousandth of the time (0.1 s) at which the NH2Br absorbance was measured with the stopped-flow system. The molar absorptivity of NH2Br was calculated by dividing the absorbance value obtained by the concentration
d
Buffer concentration not
TABLE 2. Molar Absorptivities of Bromamines at Selected Wavelengthsa compd
λ ) 232 nm
λ ) 258 nm
λ ) 278 nm
NH2Br NHBr2 NBr3
82 ( 6 2000 ( 61 3810c
273 ( 8 884 ( 27 5000c
425 ( 8 715 ( 17 1400b
a Unless otherwise indicated, the values were measured in the present study. b Value derived from data by Galal-Gorchev and Morris (3). c Value reported by Cromer et al. (13).
of HOBr originally added. The resulting molar absorptivities of NH2Br at the wavelengths of 232, 258, and 278 nm are presented in Table 2. Some discrepancies were observed between some of the molar absorptivities obtained in this study and those reported in the literature. The values reported by Galal-Gorchev and Morris (3) were 25, 265, and 390 M-1 cm-1 at the wavelengths of 232, 258, and 278 nm. However, these values were obtained by determining the bromamine concentration by iodometric titration several minutes after mixing HOBr with NH3. Consequently, based on the kinetics presented subsequently, the disproportionation of NH2Br into NHBr2 (see reaction 1) was well under way, and NHBr2 had likely interfered with the measurement (23). The molar absorptivity of NHBr2 was measured by mixing NH3 and HOBr at a [NH3]/[HOBr] ratio of 20 and pH 6.51 or pH 6.74. The reaction was allowed to proceed for at least 10 s (set ABS-2, Table 1) to ensure that NHBr2 was the only bromamine species present in solution at measurable levels. Samples were taken at various times for NHBr2 analysis by iodometric titrimetic method. Notice that the initial concentration of HOBr could not be used because even though NHBr2 was the only bromamine species present, it was decomposing continuously. The resulting concentrations and absorbances measured at the same reaction time were used to calculate the molar absorptivities presented in Table 2. In contrast to the discrepancies observed between NH2Br molar absorptivities obtained in this study and those presented in the literature, reasonably good agreement was observed for NHBr2 values. For example, the values presented in Table 2 are close to those of 1900, 900, and 700 M-1 cm-1 reported by Galal-Gorchev and Morris (3) at the respective wavelengths of 232, 258 and 278 nm or those of 900 and 780 M-1 cm-1 obtained by Cromer et al. (13) at the wavelengths of 258 and 278 nm. The molar absorptivity values obtained in this study for NH2Br and NHBr2 (Table 2) were used for the analysis of bromamines kinetic data presented in subsequent sections. VOL. 38, NO. 7, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. Bromamine decomposition at pH 6.71 and various initial bromide ion concentrations (set Br-eff, Table 1) ([HOBr]0 ) 0.55 mM, initial total ammonia ) 10 mM, 0.01 M phosphate buffer, µ ) 0.1 M (NaClO4), 25 °C). Molar absorptivities values used for NBr3 (Table 2) reported by Galal-Gorchev and Morris (3) and Cromer et al. (13) were used in preliminary analyses of absorbance traces to confirm that NBr3 was below detection under all experimental conditions used in this study. Effect of Free Bromine Production Method. The bromine stock solutions prepared by the three different methods described in the Experimental Section had different bromide and chloride ion concentrations. Experiments (Br-eff, Table 1) were performed at pH 6.71 and pH 8.48 to assess if these two ions had any effect on the kinetics of bromamine decomposition and to decide which method to use for the remaining bromamine decomposition tests. Consistent with the observations reported by Inman and Johnson (11), chloride ion had no effect on the decomposition of bromamines under any of the conditions investigated. In contrast, although bromide ion did not have a measurable effect at pH 8.48, at pH 6.71 the presence of this ion resulted in a faster rate of NH2Br disproportionation into NHBr2 and NH3. This effect is depicted in Figure 1 for the experiments performed with the free bromine solution obtained by the ozone/Br- reaction (i.e., [Br-]/[OBr-] ≈ 0), the liquid bromine hydrolysis (i.e., [Br-]/[OBr-] ≈ 1), and the liquid bromine hydrolysis with bromide addition (i.e., [Br-]/[OBr-] ≈ 4) methods. Because the presence of bromide ion was found to interfere with the bromamine kinetics, the bromine hydrolysis method was discarded. Furthermore, because no measurable differences were observed between the curves obtained with the free bromine prepared from NaOCl/Br(i.e., [Br-]/[OBr-] ≈ 0.05) (not shown) and that shown in Figure 1 for free bromine prepared by the ozone/Br- reaction method (i.e., [Br-]/[OBr-] ≈ 0) and the fact that the ozone/ Br- reaction method was more cumbersome, it was decided to produce free bromine stock solutions by the NaOCl/Brreaction method for all subsequent experiments performed in this study. Bromamine Decomposition Model for the pH Range of 6.5-9.5. In this section, bromamine decomposition reactions are obtained by assuming that they are similar to chloramines decomposition reactions (24, 25). Furthermore, simplifications are made based on the range of experimental conditions used in this study (i.e., pH range of 6.5-9.5) and ammonia to bromine molar ratios greater than 5. First, the disproportionation of monobromamine into dibromamine and ammonia (reaction 1) is consistent with a parallel reaction reported for the chloramines system (24, 25). The results obtained for experimental sets NN-1 and NN-2 were consistent with the occurrence of this reaction. For example, the data for set NN-1 (Figure 2) show that after an initial catalysis effect by ammonium ion, which will be discussed in a subsequent section, increasing the ammonia concentration resulted in a shifting to the left according to 2114
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FIGURE 2. Ammonia effect on bromamine decomposition for experiments performed at pH 8.98 (set NN-1, Table 1) ([HOBr]0 ) 0.40 mM, initial total ammonia ) 5-40 mM, µ ) 0.1 M (NaClO4), 25 °C). reaction 1 (i.e., higher NH2Br and lower NHBr2 levels). Notice that, in Figure 2 and subsequent figures describing bromamine concentrations, the scales of the y-axes for NH2Br and NHBr2 are different in order to illustrate clearly the concentrations of the two species. However, as also depicted in Figure 2, the total concentration of bromamines decreased with time as a result of decomposition. Assuming that the same reactions responsible for the decomposition of chloramines (24, 25) are applicable to the bromamines system, the following reactions would apply:
NHBr2 + H2O f I
(6)
I + NHBr2 f HOBr + products
(7)
I + NH2Br f products
(8)
k3
NH2Br + NHBr2 98 products
(9)
in which once again following the similarity with the chloramines system, I would be an unidentified intermediate compound that might absorb light in the same wavelength range as the two bromamines of interest. However, similar to the chloramine system, absorbance by the intermediate compound was not detected, indicating that its concentration remained very low, possibly resulting from reaction 6 being the rate-limiting step. Consequently, the system of reactions 6-9 could be simplified to reaction 9 and k2
2NHBr2 98 products
(10)
Notice that reaction 10 is consistent with the observation by Cromer et al. (13) that dibromamine undergoes second-order decomposition with respect to itself under conditions at which the concentration of NH2Br is low. In summary, the decomposition kinetics of bromamines can be represented with reactions 1, 9, and 10. Although based on the analogy with the chloramines system, other reactions such as dibromamine hydrolysis, and various reactions involving HOBr and tribromamine could also occur, but their effect will be neglected. This simplified approach is justified as least in part because no HOBr or NBr3 was
SCHEME 1
detected under any of the conditions investigated, and the hydrolysis of NHBr2 was assumed to be slow as compared to NHBr2 decomposition through reactions 9 and 10. Although the identification of products resulting from reactions 9 and 10 was beyond the scope of the present study, it was important to check that none of these compounds would interfere in the spectrophometric analyses of NH2Br and NHBr2. A comparison of NH2Br and NHBr2 concentrations obtained by titration and spectrophotometric methods as well as a comparison of spectra evolution over time confirmed that the reaction products did not have a measurable interference at the wavelengths of interest under the range of conditions investigated. Catalysis Effects. It was anticipated that similar to the chloramines system, bromamine reactions would also undergo catalysis (26). Experiments performed to characterize the catalysis of reaction 1 by various relevant acids as well as to assess if any catalysis also takes place for reactions 9 and 10 are presented in this section. The experiments were performed without buffer addition other than ammonia as well as with two common buffers: phosphate and carbonate (sets CN and PP). Ammonia Catalysis. Experiments performed without any buffer other than ammonia were performed at relatively high pH (Table 1) to allow some degree of pH control by the buffer capacity provided by ammonia. Ammonium ion catalyzed the conversion of NH2Br into NHBr2 (reaction 1), most noticeably at early reaction times as depicted in Figure 2. This was consistent with the occurrence of general acid catalysis (Scheme 1) similar to that reported for the disproportionation of monochloramine (27). Ammonia catalysis effects on bromamine decomposition were determined first by fitting experimental sets NN-1 and NN-2 with Dynafit (21). Each experiment within a set was fitted independently to rate expressions corresponding to reactions 1 and 9 with the rate constants k1, k-1, and k3 being used as fitting parameters. Reaction 10 was not considered because preliminary fitting efforts including all three reactions revealed that the rate of this reaction at the relatively high pH of 8.999.44 was much slower than that of competing reaction 9 and that taking it into account did not affect the fitting results. The fitted curves, shown as dashed lines, are compared to experimental data in Figure 2. As depicted in the figure, the model fitted the data well. The reaction rate constants k1 and k-1 resulting from the fitting effort are plotted in Figure 3 against ammonium ion concentration. The linearity of the plot and a comparison of the results confirmed that reaction 1 was catalyzed by ammonium ion. The ammonium ion catalysis reaction rate constants k1,NH4 and k-1,NH4 corresponding to the slopes of the linear plots in Figure 3 are presented in Table 3. The rate constant for reaction 9, k3, was found to be independent of ammonia concentration at both pH values investigated. The average k3 values were 7.32 M-1 s-1 at pH 8.98 and 9.38 M-1 s-1 at pH 9.44. The higher rate constant at higher OH- concentration was consistent with reaction 9
FIGURE 3. Dependence of fitted reaction rate constants k1 and k-1 on NH4+ concentration (sets NN-1 and NN-2). Lines represent leastsquares fit to the data (pH 8.99 and 9.44, [HOBr]0 ) 0.40 mM, total initial ammonia ) 5-40 mM, µ ) 0.1 M (NaClO4), 25 °C). being base-catalyzed (Scheme 2). The corresponding hydroxide ion and water assisted reaction rate constants, k3,OH and k3,0, are presented in Table 4. Notice that the products of the reaction in Scheme 2 are not identified. Consistent with what it has been reported for NHCl2 reactions (28, 29), it is likely that the transition state would form a N-N compound that ultimately produces N2. Hydrogen Ion Catalysis. The occurrence of catalysis by H+ was assessed with the data from experimental sets HN-1-3. These experiments were performed at pH 9.01-9.50 and initial bromamine concentrations ranging from 0.15 to 0.50 mM. Once again, the pH range was selected so that ammonia would provide the buffer capacity needed to maintain the pH constant without the addition of another buffer chemical. The total ammonia concentration (Table 1) was increased for increasing pH to achieve a target constant concentration for ammonium ion of 5.6 mM in all cases. The ammonium ion catalysis terms of the rate expression for reaction 1 (i.e., k1,NH4[NH4+] and k-1,NH4[NH4+]) were therefore kept constant. Results obtained for the experiments performed at the highest initial combined bromine concentration of 0.5 mM are shown in Figure 4. These data as well as the experimental results obtained at the other combined bromine concentrations were fitted with rate expressions corresponding to reactions 1 and 9. Once again, at the relatively high pH values used for these experiments, reaction 10 could be neglected. All experiments from sets HN-1-3 were then fitted simultaneously to the rate expressions for reactions 1 and 9. The fitting parameters were the apparent forward and reverse rate constants for reaction 1. The rate constant k3 for reaction 9 was assumed to be k3 ) k3,0 + k3,OH[OH-] using the k3,0 and k3,OH values obtained from the ammonia catalysis experiments (Table 4). Ammonia’s contribution to base catalysis of reaction 9 was neglected because it is a weaker base than OH-. The fitted curves (dashed lines) are compared to the experimental data in Figure 4. The agreement was considered to be acceptable despite the discrepancies observed for NHBr2 at relatively long reaction times. These deviations might have resulted, VOL. 38, NO. 7, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 3. Summary of Measured and Predicted Reaction Rate Constantsa for the Disproportionation of NH2Br into NHBr2 and NH3 (Reaction 1) at 25 °C and Ionic Strength of 0.1 M (NaClO4) k1,catalyst
a
catalyst
pKa (19)
p/q
H2O HPO42HCO3NH4+ H2PO4H+ H2CO3 H3PO4
15.46b
2/3 1/3 1/3 4/4 2/3 3/2 2/3 3/2
11.74 10.00 9.29 6.72 -1.72b 6.16 2.0
measured
k-1,catalyst predicted
0.5
measured 1.0
9.6 540 290 3.4 × 105 5.0 × 108
predicted
3.8 × 104 1.4 × 107
6.5 180 190 6.5 × 104 1.0 × 109
2.8 × 104 1.0 × 107
k1,catalyst and k-1,catalyst values in M-2 s-1 except the H2O terms, k1,0 and k-1,0, in M-1 s-1. b pKa values from Kumar et al. (28).
SCHEME 2
TABLE 4. Summary of Measured and Predicted Reaction Rate Constantsa for the Decomposition of NH2Br and NHBr2 (Reactions 9 and 10) at 25 °C and Ionic Strength of 0.1 M (NaClO4) k2,catalyst catalyst pKa (19) p/q OH-
15.46b
PO4311.74 CO3210.00 NH3 9.29 HPO426.72 H2O -1.72b
measured
2/3 1/3 1/3 4/4 2/3 1.4 × 103 3/2 8.9
measured
4.1 × 8.3 × 1.7 × 105 4 2.2 × 10 3.2 × 103 6.8 × 104 107
104
predicted 1.8 × 103 1.4 × 103 76
6.2 -2
a
k3,catalyst
predicted
-1
k2,catalyst and k3,catalyst values in M s except the H2O terms, k2,0 and k3,0, in M-1 s-1. b pKa values from Kumar et al. (28).
FIGURE 4. Acid catalysis effect on bromamine decomposition for experiments performed at varying pH (sets HN-1-3, Table 1) ([HOBr]0 ) 0.50 mM, [NH4+] ) 5.6 mM, µ ) 0.1 M (NaClO4), 25 °C). at least in part, from relatively small inaccuracies in the molar absorptivity values used, which would tend to affect more strongly the fitting of the species present at relatively low concentration, NHBr2, for the experiments shown in Figure 4. Resulting k1 and k-1 values are plotted against H+ 2116
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FIGURE 5. Dependence of fitted reaction rate constants k1 and k-1 on H+ concentration (sets HN-1-3). Lines represent least-squares fit to the data (pH 9.01-9.50, [HOBr]0 ) 0.15-0.50 mM, [NH4+] ) 5.6 mM, µ ) 0.1 M (NaClO4), 25 °C). concentration in Figure 5. The slopes and intercepts of these lines were used to calculate the water and hydrogen ion catalysis constants with the expressions k1 ) k1,0 + k1,NH4[NH4+] + k1,H[H+] and k-1 ) k-1,0 + k-1,NH4[NH4+] + k-1,H[H+] in which k1,NH4 and k-1,NH4 were assumed to be the values obtained from the ammonia catalysis experiments (Table 3). The resulting values for k1,0, k1,H, k-1,0, and k-1,H are presented in Table 3. The equilibrium constant for reaction 1 estimated from the k1,0 and k-1,0 values obtained in this study, K1 ) (k1/k-1) ) 0.5, was in reasonably good agreement with the value of K1 ) 0.78 estimated from free energy of formation values reported by Sugam and Helz (30) for the reactant and products of reaction 1. Carbonate Buffer Catalysis. Carbonate is the most common buffer system encountered in natural water. Experimental sets CN-1 and CN-2 were performed to investigate the occurrence of catalysis by carbonate species. The experiments were performed at pH 8.98 and pH 9.43, and so once again only reactions 1 and 9 needed to be taken into consideration. Results obtained for the experiments performed at pH 9.43 are shown in Figure 6. As depicted in the figure, the rate of NH2Br disproportionation into NHBr2 increased with increasing total carbonate concentration. The two experimental sets were fitted simultaneously to the rate expressions for reactions 1 and 9. Once again good agreement was observed between the fitted and experimental curves for NH2Br, the predominant species under the conditions investigated, while deviations can be observed for NHBr2, which was present at
FIGURE 6. Carbonate buffer effect on bromamine decomposition for experiments performed at pH 9.43 (set CN-2, Table 1) ([HOBr]0 ) 0.40 mM, total initial ammonia ) 10 mM, CT,CO3 ) 5-40 mM, µ ) 0.1 M (NaClO4), 25 °C). much lower levels. Notice that negative values were obtained for NHBr2 concentration, once again also as a result of small inaccuracies in the molar absorptivity values. The resulting k1 and k-1 values were found to increase linearly with bicarbonate ion concentration, and k3 increased linearly with increasing carbonate ion concentration. The slopes of linear plots were used for calculating the corresponding k1,HCO3, k-1,HCO3 and k3,CO3 values (Tables 3 and 4) using the following expressions:
k1,HCO3[HCO3-] ) k1 - k1,0 - k1,H[H+] - k1,NH4[NH4+] (11) k-1,HCO3[HCO3-] ) k-1 - k-1,0 - k-1,H[H+] k-1,NH4[NH4+] (12) k3,CO3[CO32-] ) k3 - k3,0 - k3,OH[OH-]
(13)
in which the k1,0, k1,H, k1,NH4, k-1,0, k-1,H, k-1,NH4, k3,0, and k3,OH are the values obtained previously from the ammonia and hydrogen ion catalysis experiments (Tables 3 and 4). Phosphate Buffer Catalysis. The decomposition of bromamines was studied at a wider pH range in phosphatebuffered solutions. The relative concentrations of bromamine during decomposition at pH below 8 are different from those observed for the experiments at relatively high pH. At the lower pH levels investigated in this study, NHBr2 was the dominant bromamine species after equilibrium between NH2Br and NHBr2 was approached. Furthermore, overall bromamine decomposition according to reaction 10 could not be neglected at the corresponding high NHBr2 concentrations. Experiments were performed at varying phosphate buffer concentration with initial ammonia and bromamine concentrations kept constant (set PP) and varying bromamine concentration with phosphate buffer and ammonia concentrations kept constant (sets HP-1-3). The pH range investigated with all sets was 6.77-8.81 (Table 1). Results obtained for experimental set PP are shown in Figure 7. As depicted in the plots, the rate of disproportionation of NH2Br into NHBr2 and NH3 increased with increasing total phosphate concentration. The shape of the
FIGURE 7. Phosphate buffer effect on bromamine decomposition for experiments performed at pH 7.21 (set PP, Table 1) ([HOBr]0 ) 0.50 mM, total initial ammonia ) 10 mM, CT,PO4 ) 5-40 mM, µ ) 0.1 M (NaClO4), 25 °C). curves revealed that a relatively fast equilibration between the bromamine species consistent with reaction 1 was followed by slower bromamine decay. The effect of phosphate buffer on the initial equilibration is consistent with the occurrence of acid catalysis induced by H2PO4-. Although not as noticeably in Figure 7, the rate of bromamine decomposition also increased with increasing phosphate concentration. Each experiment of set PP was fitted with the rate expressions for reactions 1, 9, and 10 to obtain the corresponding rate constants k1, k-1, and k2. The rate constant k3 was represented with the expression k3 ) k3,0 + k3,OH[OH-] using the k3,0 and k3,OH values obtained from the ammonia catalysis experiments (Table 4). As shown in Figure 7, the fitted curves (dashed lines) were in good agreement with the experimental data. The phosphate catalysis rate constants were then calculated from the fitting parameters k1, k-1, and k2. The phosphate catalysis terms for reaction 1 were obtained with the following expressions:
k1,H2PO4[H2PO4-] ) k1 - k1,0 - k1,H[H+] - k1,NH4[NH4+] (14) k-1,H2PO4[H2PO4-] ) k-1 - k-1,0 - k-1,H[H+] k-1,NH4[NH4+] (15) in which k1,0, k1,H, k1,NH4, k-1,0, k-1,H, and k-1,NH4 are the values obtained from the ammonia and hydrogen ion catalysis experiments (Table 3). The phosphate catalysis terms calculated with equations 14 and 15 are plotted against the corresponding concentration of H2PO4- in Figure 8. The linear relationships depicted in the plot confirmed the occurrence of acid catalysis of reaction 1 by H2PO4-. The slope of the regression lines shown in Figure 8 were used to obtain the k1,H2PO4 and k-1,H2PO4 values listed in Table 3. A linear plot of fitting parameter k2 against the corresponding concentration of HPO42- is shown in Figure 9. Once again a linear plot was obtained consistent with the occurrence of base catalysis by HPO42-. The slope and intercept of the plot were used to calculate the k2,0 and k2,HPO4 values listed in Table 4. These results confirmed that, similar to the chloramines system, the disproportionation of NH2Br into NHBr2 and NH3 (reaction 1) is general acid catalyzed in both directions, and the subsequent decomposition of both VOL. 38, NO. 7, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 8. Dependence of fitted reaction rate constants ki ) k1k1,0-k1,H[H+]-k1,NH4[NH4+] and k-i ) k-1-k-1,0-k-1,H[H+]-k-1,NH4[NH4+], on H2PO4- concentration (sets PP and HP-1-3). Lines represent least-squares fit to the data (pH 6.77-8.81, µ ) 0.1 M (NaClO4), 25 °C).
FIGURE 10. Bromamine kinetics at various ammonia concentrations (set NP, Table 1). Symbols are experimental results, and lines are model predictions (pH 7.95, [HOBr]0 ) 0.40 mM, 0.01 M phosphate buffer, µ ) 0.1 M (NaClO4), 25 °C). FIGURE 9. Dependence of fitted reaction rate constant k2 on HPO42concentration (sets PP and HP-1-3). Lines represent least-squares fit to the data (pH 6.77-8.81, µ ) 0.1 M (NaClO4), 25 °C). bromamine species (reactions 9 and 10) are general base catalyzed, although not all catalysis terms could be determined at the range of experimental conditions used. Prediction of Bromamine Decomposition. Bromamine decomposition in phophate buffer with varying ammonia concentrations (set NP, Table 1) was predicted with the rate expressions for reactions 1, 9, and 10 using the following expressions for the corresponding rate constants:
k1 ) k1,0 + k1,H[H+] + k1,NH4[NH4+] + k1,H2PO4[H2PO4-] (16) k-1 ) k-1,0 + k-1,H[H+] + k-1,NH4[NH4+] + k-1,H2PO4[H2PO4-] (17) k2 ) k2,0 + k2,HPO4[HPO42-]
(18)
k3 ) k3,0 + k3,OH[OH-]
(19)
with the rate constant values presented in Tables 3 and 4. The predicted curves for NH2Br and NHBr2 concentrations as a function of reaction time are compared to experimental data in Figure 10. As depicted in the figure, though some deviations were observed at the highest ammonia concentrations investigated, good general agreement was observed between predictions and data. Brønsted Relationship. The various catalysis rate constants summarized in Tables 3 and 4 were further used to check if they were correlated with the acid dissociation equilibrium constants and consistent with the Brønsted relationship (31):
()
log 2118
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( )
Kaq k ) log GA + R log p p
(20)
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FIGURE 11. Brønsted plot for catalysis terms of rate constant k1 (reaction 1). Solid symbols are values obtained experimentally, and open symbols are values calculated with the regression line.
in which p is the maximum number of protons that can be donated by the acid form and q is the maximum number of protons that can be accepted by the conjugate base. The parameter R indicates the degree of proton-transfer ability in the transition state. The p and q parameters for each acidbase system of interest in this study are summarized in Table 3. As depicted by the resulting plots in Figures 11-13, the rate constants obtained were generally consistent with equation 20. The slopes of the plots, ranging from R ) 0.51 to R ) 0.63 were indicative of relatively moderate degree of proton transfer from an acid (HA in Scheme 1) to NH2Br or from NHBr2 to a base (B- in Scheme 2). The resulting correlations could now be used to estimate the contribution by other relevant acids or bases to the overall acid catalysis of reaction 1 and base catalysis of reactions 9 and 10. The predicted rate constant values are presented in Tables 3 and 4.
FIGURE 12. Brønsted plot for catalysis terms of rate constant k-1 (reaction 1). Solid symbols are values obtained experimentally, and open symbols are values calculated with the regression line.
FIGURE 13. Brønsted plots for catalysis terms of rate constants k2 (reaction 10) and k3 (reaction 9). Solid symbols are values obtained experimentally, and open symbols are values calculated with the corresponding regression lines.
Acknowledgments This research was supported by a grant from the U.S. Environmental Protection Agency Science to Achieve Results (STAR) Program (Grant R826830-01). The authors would like to thank Dr. Urs von Gunten from EAWAG (Duebendorf, Switzerland) and Dr. Chad Jafvert from the School of Civil and Environmental Engineering, Purdue University (West Lafayette, IN) for fruitful discussions.
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Received for review July 7, 2003. Revised manuscript received January 12, 2004. Accepted January 13, 2004. ES034726H
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