- Email: [email protected]
Rate constants of reactions of bromine with phenols in aqueous solution
Water Research 37 (2003) 2883–2892
Rate constants of reactions of bromine with phenols in aqueous solution Herve! Gallard*, Fabien Pellizzari, Jean Philippe Croue! , B. Legube Laboratoire de Chimie de l’Eau et de l’Environnement, UMR CNRS 6008, Ecole Sup!erieure d’Ing!enieurs de Poitiers 40, Avenue du Recteur Pineau, Poitiers Cedex 86 022, France Received 15 May 2002; accepted 20 February 2003
Abstract The kinetics of bromination of six ortho- and para-substituted phenols was investigated between pH 5 and pH 12 in aqueous solution. Kinetics was followed with a continuous-flow reactor previously validated by studying the fast reaction between chlorine and ammonia. The overall reaction rate between bromine and phenols is controlled by the reaction of HOBr with the phenoxide ion between pH 6 and pH 10. The reaction of HOBr with the undissociated phenols and the reaction of BrO with the phenoxide ions become only significant for pHo6 and pH>10, respectively. The second-order rate constants for the reaction of HOBr with phenoxide ions vary between 1.4(70.1) 103 and 2.1(70.5) 108 M1 s1 for 2,4,6-trichlorophenol and 4-methylphenol, respectively. Hammett-type correlation was obtained for the reaction of HOBr with the phenoxide ions (log(k)=8.03.33 Ss) and was compared with Hammett-type correlations of HOCl and HOI. The reaction rate of bromine with phenol-like organic compounds was estimated to be about 103-fold higher than with chlorine and 103-fold lower than with ozone in drinking water treatment conditions. r 2003 Elsevier Science Ltd. All rights reserved. Keywords: Bromination; Phenols; Rate constants; Disinfection; Drinking water
1. Introduction In drinking water treatment plants, hypobromous acid (HOBr) is formed as a disinfectant by-product when chlorine or ozone is used to treat bromidecontaining natural waters. Once formed, hypobromous acid dissociates into hypobromite ion (BrO) with a pKa of 8.89: HOBr"BrO þ Hþ ;
Ka :
ð1Þ
*Corresponding author. Tel.: +33-5-49-45-44-31; fax: +335-49-45-97-68. E-mail address: [email protected] (H. Gallard).
In the presence of ammonia hypobromous acid reacts in a fast reaction producing monobromamine (NH2Br) and dibromamine (NHBr2) [1]. Hypobromous acid reacts also with natural organic matter (NOM) in complex reactions leading to the formation of bromoorganic compounds. The formation of brominated and mixed chlorobromo by-products such as bromoform (CHBr3) and mixed trihalomethanes (THMs) (dichlorobromomethane (CHCl2Br), dibromochloromethane (CHBr2Cl)) have been reported in finished drinking waters since the early 1970s [2,3]. Toxicological studies conducted on laboratory animals gave the first findings linking brominated by-products to colon cancer [4]. According to Morris [5], 5000 cases of bladder cancer per year and 8000 cases of rectal cancer per year may be associated with the
0043-1354/03/$ - see front matter r 2003 Elsevier Science Ltd. All rights reserved. doi:10.1016/S0043-1354(03)00132-5
2884
H. Gallard et al. / Water Research 37 (2003) 2883–2892
consumption of chlorinated drinking water in the US. The interest on trihalomethanes increased recently when epidemiological studies showed increased risks for birth defects, low birth weight, and miscarriages associated with the presence of THMs in drinking waters [6]. In the ozonation step, hypobromous acid reacts with ozone and OH radical. Hypobromous acid is the key intermediate in the formation mechanism of bromate (BrO 3 ), a potential carcinogenic by-product of ozonation in waterworks [7,8]. Kinetics of reaction of HOBr with ammonia, ozone, H2O2, and radicals has been extensively studied. The rate constants are now rather well known and the formation of bromate has been modeled in synthetic water [8]. In natural waters, the task appears more difficult because NOM can act as a scavenger of HOBr. To understand and evaluate bromate formation in natural waters, it was then concluded that it is necessary to better understand the NOM reactions with oxy-bromine species [9]. Number of studies has shown that the presence of bromide during chlorination increases the yield and the rate of formation of THMs [10,11,2]. To better understand the reactivity of chlorine or bromine with NOM, chlorination and, to a lesser extent, bromination of model organic compounds of NOM, such as phenolic compounds, have been studied. Bromination of resorcinol and orcinol yielded mono-, di- and tribrominated electrophilic substitution products in meta and para positions [12]. It is expected that 1,3-dihydroxy aromatic substrates would react with chlorine and bromine via the same pathway leading to the formation of chloroform and bromoform, respectively. The kinetics of bromination of phenols was studied by Bell and Rawlinson [13] and Tee et al. [14] at pHo7.0 in the presence of high concentration of bromide (2.5 mMo[KBr]o1.5 M). In such conditions, dibromine (Br2) and tribromide ion (Br 3 ) were considered as the only brominating species. Tee et al. [14] obtained, for phenol, a value of 1.2 109 M1 s1 for the attack of bromine on phenoxide ion and 8.5 108 M1 s1 for the reaction with tribromide ion. In the absence of bromide, the rate of bromination of aromatic compounds is proportional to the acidity of the media at low concentrations of acids [15]. It was hypothesized that the rate-limiting step is either the reaction of the positive bromine species H2OBr+ with organic compounds or the decomposition of a protonated transition state formed between HOBr and the organic substrate. An apparent second-order rate constant k ¼ 500 M1 s1 for the reaction between HOBr and phenol was determined in aqueous solution at pH 4.0 in batch reactor [16]. However, it is important to point out that for all these studies, pHs are too
low or the concentrations of bromide are too high to describe the reactivity of bromine with organic compounds in drinking water treatment. The kinetics of reaction of HOI and HOCl with phenols was studied in aqueous solutions between pH 2 and pH 12 [17,18]. The second-order rate constants of the reaction showed a pH-dependence that was explained by the speciation of both halogen and phenol species. For chlorine, apparent rate constants showed a maximum in the pH region 8–9, which is explained if HOCl is the only active electrophile and the reactivity of OCl is negligible. For hypobromous acid, a similar behavior is expected with a maximum shifted to higher pH because of a higher pKa for bromine. Few data are available regarding the kinetics of bromine in aqueous solution at near neutral pH in the absence of bromide. The purpose of this study was to determine the rate constants of initial bromination of selected phenolic compounds between pH 5 and pH 12.
2. Experimental section All chemicals were of the purest available quality. Solutions of organic compounds, ammonia, chlorine and bromine were prepared with ultra-pure water produced from a MilliQ (Millipore) water purification system. A total of six phenolic compounds was studied: phenol, p-cresol, p-chlorophenol, p-acetylphenol, 2,4dichlorophenol and 2,4,6-trichlorophenol. Stock solutions of chlorine were prepared by diluting a commercial solution of sodium hypochlorite (NaOCl, 4% active chlorine, Aldrich). Sodium hypochlorite was standardized by iodometry. Stock solutions of bromine were prepared from a solution of ozone (5 C, pH 4.0, 10 mM phosphate) by addition of 0.1 M solution of potassium bromide (KBr) [19]. After a day, ozone was purged by bubbling nitrogen for 30 min. The bromine solution was standardized by direct photometric determination of the hypobromite ion (BrO) at 329 nm (e ¼ 332 M1 cm1 [20]) after adjusting the pH of the solution to 11. For a ratio Br/O3=0.9, the yield of the reaction was X95% and the concentration of the bromine solution was typically 0.65 mM. The solution was stable for 4 weeks when stored at 4 C. All experiments were performed at room temperature (2372 C). pH measurements were carried out with a Tacussel PHM210 pH meter, which was calibrated with pH 4, 7 and 9 standard buffers (Merck). Solutions were buffered using phosphate (5 mM). Spectrophotometric measurements were performed on a SAFAS 190 DES spectrophotometer (SAFAS, Monaco).
H. Gallard et al. / Water Research 37 (2003) 2883–2892
A preliminary study showed that the initial reaction of bromine with most of the phenolic compounds used in this study was too fast to be followed in a batch reactor. Therefore, the kinetics of bromination was studied using both batch reactor and continuous-flow reactor.
3. Batch reactor experiments The batch reactor consisted of a 1-L glass bottle equipped with a dispenser. Experiments were initiated by adding an aliquot of a stock solution of 2,4,6trichlorophenol with micropipettes to the brominecontaining aqueous solution. At different reaction times, 3 mL of solution was rapidly transferred into a vial containing 0.1 mL of sodium sulfite to quench the reaction. Concentrations of 2,4,6-trichlorophenol were determined by HPLC (column ResolvTM 5 mm Spherical C18 waters; flow rate 1 mL/min; injection volume 200 mL; mobile phase 60% methanol/40% water/0.1% acetic acid; UV detection at 215 nm).
5.1. Validation of the continuous-flow apparatus The confirmation of the kinetics of the reaction between chlorine and ammonia was used to validate the measurements made using the continuous-flow apparatus. The reaction between hypochlorous acid (HOCl) and ammonia (NH3) is a fast, second-order reaction (reaction (2))—first-order with respect to chlorine and first-order with respect to ammonia. At 20 C, the reaction rate constant, kcl ; between the two neutral molecules is 5.6 106 M1 s1 [22]: HOCl þ NH3 -products; kcl :
For the other five selected phenolic compounds the experimental setup consisted of two HPLC pumps (Constametrics3200 LDC Analytical). The first one was used to pump the solution of oxidant (bromine or chlorine), the second one the solution of the tested molecule (phenolic compounds or ammonia). Both solutions were pumped with the same flow rate. Total flow rates in the reactor varied between 2 and 18 mL min1. The two solutions were mixed using a 1/ 1600 mixing tee, connected to a 78.0 cm or a 28.1 cm reaction tube (0.25-mm-ID Peek tube) where the reaction took place. The reaction solution was collected in a 5-cm long photometric cell containing 3 mL of a buffered (pH 6.4) solution of DPD. The exact collected volume was determined using a precision balance (70.1 mg). The DPD solution stopped the reaction and allowed analysis of the residual bromine (or chlorine) in solution at 510 nm [21]. Detection limit was 0.09 mM. The reaction time was defined based on the volume of the tube and the flux of the mixed solution. Reaction times ranged about from 10 to 150 ms.
5. Results and discussion A preliminary study was conducted to test the suitability of the continuous-flow apparatus for the study of fast kinetics.
ð2Þ
Both concentrations of HOCl and NH3 depend on pH according to equilibria (3) and (4): HOCl"ClO þ Hþ ; þ NHþ 4 "NH3 þ H ;
pKa2 ¼ 7:54; pKa3 ¼ 9:3:
ð3Þ ð4Þ
Considering [Cl2]t=[HOCl]+[OCl ] and [NH3]t= [NH3]+[NH+ 4 ] and both equilibria (3) and (4), the overall rate of chlorine consumption can be described as follows:
4. Continuous-flow apparatus experiments
2885
d½Cl2 t ¼ kapp ½Cl2 t ½NH3 t ; dt
ð5Þ
where kapp ¼ kcl ; ½Hþ
Ka3 : ½H þ Ka2 ½Hþ þ Ka3 þ
ð6Þ
The apparent reaction-rate constant, kapp ; based on the total concentration of chlorine and ammonia depends on the individual reaction-rate constant kcl ; pH and the equilibrium constants Ka2 and Ka3 (Eq. (6)). The experimental pH profile of kapp was determined between pH 5.6 and pH 10.6 under pseudo-first-order kinetic experimental conditions ½NH3 t b½Cl2 t : The initial concentration of ammonia varied between 0.1 and 2.4 mM. The initial concentration of chlorine was 5 mM. Inset in Fig. 1 shows, for three different experimental conditions, the plot of Ln ½Cl2 t =½Cl2 t0 vs. time (to150 ms). The observed pseudo-first-order rate constant, kobs ; was determined as the slope of the linear curve: ½Cl2 t Ln ¼ kobs t; ð7Þ ½Cl2 t0 where the observed rate constant kobs ¼ kapp ½NH3 t0 :
ð8Þ
Fig. 1 depicts the experimental and theoretical profile of the apparent second-order rate constant (kapp ) vs. pH. The theoretical pH profile was determined by a non-linear least-squares regression of experimental data using Eq. (6) and the fit curve function of SigmaPlot 2000 software as a tool. The calculated value of kcl was 5.97(70.54) 106 M1 s1. This value ob-
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2886
4 Ln ([Cl2]t /[Cl2]t0)
-1
Log (kapp) (M s-1)
5
3
0 -0.2 -0.4 -0.6 -0.8 -1 -1.2
pH = 7.1 pH = 7.4
0
0.05
pH = 8.9 0.1
0.15
Time (sec.)
2 5
6
7
8 pH
9
10
11
Fig. 1. Apparent rate constant pH profile and pseudo-firstorder kinetic plot (inset) of the reaction of chlorine with ammonia. Conditions: 23(72) C, 0.1 mMo[NH3]t0o2.4 mM, [Cl2]t0=5 mM.
tained with the continuous-flow apparatus is close to the value of 5.6 106 M1 s1 given by Saguinsin and Morris [22]. It was concluded that this experimental setup can be conveniently used for studying the fast reactions of halogens, e.g. the reaction of bromine with phenols. 5.2. Determination of reaction rate constants between bromine and highly reactive phenolic species Based on the kinetic study of iodination [17] and chlorination [23,18] at neutral and alkaline pH, the initial bromination of phenolic compounds was considered as a second-order reaction (reaction (9))—firstorder in active bromine ([Br2]t=[HOBr]+[BrO]) and first-order in phenol ([Ph]t=[PhOH]+[PhO]): Ph þ Br2 -Products; k:
ð9Þ
Considering the speciation of both bromine (reaction (1), pKa ¼ 8:89) and phenolic compound (reaction (10)), the second-order rate constant, k; based on the total concentration of active bromine and phenolic compound, will depend on the degree of dissociation of all species: PhOH$PhO þ Hþ ;
Ka4 :
ð10Þ
Between pH 5 and pH 12, hypobromous acid is probably the most reactive species and can react with both undissociated and dissociated phenolic species according to reactions (11) and (12), respectively. The hypobromite ion is less reactive and might only react with phenoxide ion in alkaline solutions (reaction (13)): HOBr þ PhOH-Products; k1 ;
ð11Þ
HOBr þ PhO -Products;
k2 ;
ð12Þ
BrO þ PhO -Products;
k3 :
ð13Þ
In the first assumption, the dibromine (Br2) species was not introduced into the model. Considering the yield of 95% for the conversion of Br into active bromine with ozone (see experimental section) and an equilibrium constant of 5.8 109 for reaction (14), the proportion of dibromine was calculated to be o0:08% at pH 5. Its reaction with phenol was then considered negligible because the kinetic constants of HOBr reported later in the paper are only one order of magnitude smaller than the Br2 rate constant determined by Tee et al. [14] for the same compound; results that reinforce our hypothesis: HOBr þ Br "Br2 þ OH :
ð14Þ
The reaction between OBr and PhOH was considered to be negligible, because when introduced into the model, the best fit was obtained when the k value tends to zero. Also, one should note that in the literature the reactions between OX and phenol species are generally not included into models. Our results showed that the contribution of the OBr/PhO reaction seems to be significant only above pH 11. Consequently, the consumption of active bromine can be written as follows: d½Br2 t ¼ k1 ½HOBr ½PhOH k2 ½HOBr ½PhO
dt k3 ½BrO ½PhO ;
ð15Þ
d½Br2 t ¼ ððk1 ð1 aBr Þð1 aPh Þ þ k2 ð1 aBr ÞaPh dt þ k3 aBr aPh ÞÞ½Br2 t ½Ph t ; ð16Þ kapp ¼ k1 ð1 aBr Þð1 aPh Þ þ k2 ð1 aBr ÞaPh þ k3 aBr aPh ; ð17Þ where aBr and aPh are the degrees of dissociation of hypobromous acid (aBr =[BrO]/([HOBr]+[BrO]) and phenolic compound (aPh =[PhO]/([PhO]+[PhOH])). The degrees of dissociation can be calculated from the pH and the dissociation constants as follows: aBr ¼
Ka1 ; ðKa1 þ ½Hþ Þ
ð18Þ
aPh ¼
Ka4 : ðKa4 þ ½Hþ Þ
ð19Þ
Replacing aBr and aPh by Eqs. (18) and (19) into Eq. (17), the second-order rate constant, kapp ; depend on the individual reaction-rate constants (k1 ; k2 ; k3 ), the pH, and the acid dissociation constants according to
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2887
Table 1 pKa ; Hammett constants and rate constants of HOBr (k1 ; k2 ) and OBr (k3 ) with selected phenolic compounds Phenolic compounds
pKa
Ss
k1 (M1 s1)
k2 (M1 s1)
k3 (M1 s1)
p-Cresol Phenol 4-Chlorophenol 4-Acetylphenol 2,4-Dichlorophenol 2,4,6-Trichlorophenol
10.26 10.0 9.43 8.6 7.85 6.17
0.13 0 0.24 0.47 0.60 0.97
o1.0 104 o500 6.0(73.0) 103 1.0(70.4) 104 3.0(71.0) 104 —
2.1(70.5) 108 1.8(70.2) 108 7.0(70.8) 106 4.1(70.6) 106 8.8(70.9) 105 1.4(70.1) 103
o3.0 105 o1.0 105 5.5(74.0) 104 o2.0 103 3.0(72.0) 104 p1
0
6 pH 12.0 -1 -1
Log (kapp) en M s
-1 pH 8.9
-1.5
pH 10.1
-2
150
kobs (s -1)
Ln ([Br2]t /[Br2]to)
-0.5
-2.5
100 50
5
0
-3
0
20
40
60
[dichlorophenol]to (µM)
4.5
-3.5 0
0.05
0.1
0.15
Time (sec.)
Fig. 2. Pseudo-first-order kinetic plots of the bromination of 2,4-dichlorophenol. Conditions: 23(72) C, [2,4-dichlorophenol]t0=125 mM, [Br2]t0=8 mM. The inset shows the evolution of kobs with the initial concentration of 2,4-dichlorophenol at pH 7.9.
Eq. (20): þ 2
kapp ¼
5.5
þ
k1 ½H þ k2 ½H Ka4 þ k3 Ka1 Ka4 : ðKa1 þ ½Hþ ÞðKa4 þ ½Hþ Þ
ð20Þ
The individual rate constants were determined by a nonlinear least-squares regression of the experimental pH profile of kapp : The kinetics of initial bromination of phenol, 4chloro, 2,4-dichlorophenol, 4-acetylphenol and p-cresol were studied between pH 5 and pH 12 with the continuous-flow apparatus under experimental conditions developed for pseudo-first-order kinetics ð½Br2 t0 515 ½Ph t0 Þ: The acid dissociation constants of the studied phenols are given in Table 1. The initial concentrations of phenols varied from 25 to 125 mM. The initial concentrations of bromine ranged from 1.3 to 7.9 mM. In the present paper, the bromination of 2,4-dichlorophenol is described in detail and the results for the other selected phenols are presented less extensively.
4
5
6
7
8
9
10
11
12
13
pH Fig. 3. pH profile of the second-order apparent rate constant of initial bromination of 2,4-dichlorophenol. Conditions: 23(72) C, 25 mMo[2,4-dichlorophenol]t0o125 mM, 1.3 mMo [Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
Fig. 2 plots for three different pH values the evolution of Ln ½Br2 t =½Br2 t0 vs. time for the bromination of a 125 mM solution of 2,4-dichlorophenol. Results show that the rate of bromination can be well linearized using a pseudo-first-order kinetic model. The reaction order with respect to 2,4dichlorophenol was examined at pH 7.9 by varying the concentration of 2,4-dichlorophenol. A linear plot could be drawn between the pseudo-first-order rate constant and the initial concentration of 2,4dichlorophenol (inset in Fig. 2). Therefore, the rate of the bromine disappearance is first order in the concentrations of bromine and 2,4-dichlorophenol. Fig. 3 shows the experimental and theoretical pH profiles of kapp for the bromination of 2,4-dichlorophenol between pH 5 and pH 12. As for chlorine [18,23] and iodine [17], pH profiles exhibit a maximum in near-alkaline media, which is explained if HOBr is the major electrophilic species. Using the acid dissociation
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2888 7
-1
-1
Log (kapp) (M s )
6.5 6 p -methylphenol
5.5 5
p -chlorophenol
4.5 4 3.5 3 4
5
6
7
8 9 pH
10
11
12
13
Fig. 4. pH profile of the second-order apparent rate constant of initial bromination of p-methylphenol and p-chlorophenol. Conditions: 23(72) C, 25 mMo[Ph]t0o125 mM, 1.3 mMo [Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
7 6.5
phenol
-1
-1
Log (kapp) (M s )
6 5.5 5 4.5 p -acetylphenol
4 3.5
the pH profile of kapp obtained in our study for phenol is in good agreement with the apparent second-order rate constant found by Pinkernell and von Gunten [16] at pH 4.0 in a batch reactor (kapp ¼ 500 M1 s1). The individual reaction-rate constants of phenol and substituted phenols studied with the continuous-flow apparatus are listed in Table 1. The rate constants of HOBr with the phenoxide ions range from 2.1 108 M1 s1 for p-cresol to 8.8 105 M1 s1 for 2,4-dichlorophenol. For phenol, the rate constant of HOBr with phenoxide ion is 1.8 108 M1 s1, which is about one order of magnitude smaller than the corresponding rate constant for the Br2 species (k ¼ 1:2 109 M1 s1 [14]). This verified that, in our condition, the rate constant with dibromine is not enough to compensate the low Br2 concentrations. In agreement with the electrophilic character of HOBr the rate constants of HOBr with the phenoxide ions decrease with the deactivating effect of the substituents. These rate constants are one to four orders of magnitude greater than the corresponding rate constants of the reaction between HOBr and the undissociated phenols and the rate constants of the reaction between BrO with the phenoxide ions. In pH region 6–9, the reaction between HOBr and the phenoxide ions controls the overall reaction. As a consequence, the precision for k1 and k3 is very low and only upper limits were determined for phenol, pcresol and p-acetylphenol. 5.3. Bromination of 2,4,6-trichlorophenol
3 2.5 2 4
5
6
7
8
9
10
11
12
13
pH
Fig. 5. pH profile of the second-order apparent rate constant of initial bromination of phenol and p-acetylphenol (full and open circles: this study; open square: from [16]). Conditions: 23(72) C, 25 mMo[Ph]t0o125 mM, 1.3 mMo[Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
constants Ka1 ¼ 108:9 and Ka4 ¼ 107:85 the individual rate constants were calculated as k1 ¼ 3:0ð71:0Þ 104 M1 s1, k2 ¼ 8:8ð70:9Þ 105 M1 s1 and k3 ¼ 3:0ð72:0Þ 104 M1 s1. The same approach was used to model the bromination kinetics of the other selected phenolic compounds. Figs. 4 and 5 show the pH dependency of the secondorder rate constant for both 4-methylphenol and 4chlorophenol (Fig. 4) and for both phenol and 4acetylphenol (Fig. 5). Again, the proposed model fits well with the experimental profile. Fig. 5 shows also that
Since 2,4,6-trichlorophenol exhibits a much lower reactivity toward active bromine, kinetics of bromination were studied in a classical batch reactor under pseudo-first-order experimental conditions ([2,4,6-trichlorophenol]t05[Br2]t0). The initial concentration of 2,4,6-trichlorophenol was 0.5 or 1 mM. The concentration of active bromine was >20 mM. Experiments were conducted between pH 5.0 and pH 11.2. Between pH 7.0 and pH 11.2 the pseudo-firstorder rate constant determined from the linear plots ln[trichlorophenol]t/[trichlorophenol]t0 vs. time was found to be directly proportional to the initial concentration of 2,4,6-trichlorophenol (results not presented). The disappearance of 2,4,6-trichlorophenol is first order with respect to bromine and 2,4,6-trichlorophenol concentrations. A rate constant of 1400(7100) M1 s1 was found for the reaction of HOBr with the 2,4,6trichlorophenoxide ion by fitting experimental values with the above proposed model. Both experimental and theoretical pH profiles are plotted in Fig. 6. In this pH region, the overall reaction is only controlled by the reaction of HOBr with the 2,4,6-trichlorophenoxide ion. The rate constant of the reaction between BrO and 2,4,6-trichlorophenolate was estimated to be below
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2889
10 3
Log (k2) (M s-1)
2 -1
-1
-1
Log (kapp) (M s )
9
1 0
p- Me
Ph
8 7
HOI
y = -3.33 (±0.45) . Σσ + 7.98 2 R = 0.95
p- Cl
p -Ac
6
2,4-diCl
5
HOCl
4
-1
2,4,6-triCl
3 -2 4
6
8 pH
10
12
Fig. 6. pH profile of the second-order apparent rate constant of initial bromination (( ) this study) and chlorination [26] of 2,4,6-trichlorophenol. Conditions: 23(72) C, 25 mMo [Ph]t0o125 mM, 1.3 mMo[Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k2 and k3 to the experimental data (symbols).
1 M1 s1. For comparison the pH profile of the secondorder rate constant of chlorination of 2,4,6-trichlorophenol is also plotted in Fig. 6. Results show that the rate constants of bromination are more than two orders of magnitude greater than the rate constants of chlorination. As expected, the maximum of the pH profile is shifted towards higher pH values from chlorine to bromine due to the higher pKa of HOBr/BrO. For pHo7 and in our experimental conditions, the reaction rate between bromine and 2,4,6-trichlorophenol deviated from a second-order kinetic model. At pH 5.8, the degradation rate of 2,4,6-trichlorophenol could be linearized using a pseudo-first-order kinetic model; however, the degradation rate did not depend on the concentration of bromine. The rate of trichlorophenol disappearance was found to be zero order with respect to [HOBr]. In such conditions, the reaction rate of HOBr with the trichlorophenoxide ion is not the ratelimiting step. The rate-limiting step is the deprotonation of 2,4,6-trichlorophenol and the observed pseudo-first0 order rate constant, kobs ; is equal to the rate constant of deprotonation kd : A first-order rate constant 0 kd ¼ kobs ¼ 2:3 103 s1 was found. The rate constant of the reaction between the trichlorophenoxide ion and H+ can then be determined from the acid dissociation constant Ka4 : kf ¼ kd =Ka4 : With pKa4 ¼ 6:17 (Ka4 ¼ 6:8 107 ), a value kf ¼ 3400 M1 s1 was calculated for the reaction between 2,4,6-trichlorophenol and H+. 5.4. Linear free-energy relationships The Hammett correlation is classically used to predict the effect of substituents on reaction-rate constants
2 -0.5
0
0.5 Σσ
1
1.5
Fig. 7. Hammett plots for the reaction of HOBr, HOI and HOCl with phenoxide ions (HOBr: this study; HOCl: from [18] HOI: from [17]). p-Me: p-methylphenol; Ph: phenol; p-Cl: pchlorophenol; p-Ac: p-acetylphenol; 2,4-diCl: 2,4-dichlorophenol; and 2,4,6-triCl: 2,4,6-trichlorophenol.
between organic compounds and oxidants [23,17,18]. The Hammett constant (s) reflects the effects of substituents on the electron density of the aromatic ring by inductive and resonance effects. For s > 0; substituents have an electron-withdrawing effect (e.g. CN), whereas for so0 substituents exert an electron-donating effect (e.g. CH3). The non-substituted phenol was used as a reference compound (Sso;m;p ¼ 0). Values of s were taken from Hansch et al. [24] and the values of Sso;m;p are shown in Table 1 for the studied phenolic compounds. The Hammett correlation was only tested for the reaction between HOBr and phenoxide ions because these rate constants were determined with the best precision and also because this particular reaction controls the reactivity of bromine in natural waters (neutral pH). Fig. 7 shows that a good linear plot is obtained for log(k2 ) vs. Sso;m;p with our results. The linear regression for the reaction of bromine with the five selected phenolic compounds substituted in para or/ and in ortho position is X logðk2 Þ ¼ 8:0 3:33ð70:45Þ so;m;p ; r ¼ 0:946;
n ¼ 5:
ð21Þ
The negative sign of the Hammett slope (r) is typical of electrophilic substitution reaction because the increase of the Hammett constant corresponds to an increase of the deactivating effect of the substituents. With phenolic compounds the OH and O substituents act as activating substituents that promote the halogen substitution onto the ortho and para positions. The initial reaction of bromine
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2890
Table 2 Rate constants and pseudo-first-order reaction rates for the reaction of HOBr, HOCl and O3 with phenol and 4-chlorophenol at pH 8.0 (2372 C) Phenolic compound
Phenol 4-Chlorophenol
HOCl
HOBr
O3
k (M1 s1)b
k0 (s1)a
k (M1 s1)c
k0 (s1)a
k (M1 s1)d
k0 (s1)a
60 30
1.2 103 6.0 104
2.2 106 2.3 105
2.2 0.23
18 106 34 106
360 680
a
Pseudo-first-order rate constants are calculated from second-order rate constants for [HOCl]=20 mM, [HOBr]=1 mM and [O3]=20 mM. b From [18]. c Our study. d From [25].
with the studied phenols is therefore the electrophilic substitution onto available 2, 4 and 6 ring positions leading to the formation of bromophenols [14]. As a consequence the above correlation cannot be used to predict the rate constant of bromine with 2,4,6trichlorophenoxide ion since, in this case, all para and ortho positions are already substituted with chlorine atoms. This is confirmed in Fig. 7 in which rate constant of 2,4,6-trichlorophenoxide ion is significantly below the Hammett correlation. Fig. 7 gives for comparison the Hammett-type correlations established for the rate constants of HOCl [18] and HOI [17] with phenoxide ions. Results show that HOBr is the most reactive oxy-halides. Rate constants of HOBr are about three to four orders of magnitude higher than for HOCl and between 5 and 500-fold higher than for HOI. The intermediate reactivity of iodine is not consistent with the order of electronegativity of halogens. Because iodine has the lowest electronegativity, hypoiodous acid should present the highest rate constants. These results remain unclear. One explanation may be the low stability of HOI in aqueous solution, which makes the determination of its rate constants with organic compounds more difficult. Steric effect might also explain the unexpected low reactivity of iodine. The Hammett slopes of HOCl (r ¼ 3:0) and HOBr (r ¼ 3:33) are rather similar but the rate constants of HOI show a stronger effect of substituents with r ¼ 5:4: As it can be expected from electronegativity scale, the jrj value increases with decreasing electronegativity of halogens.
of Br). Corresponding pseudo-first-order rate constants were calculated at pH 8.0 for ozone ([O3]=20 mM; D1 mg/L) and chlorine ([HOCl]=20 mM; D1.4 mg/L) from literature data. Table 2 presents the apparent second-order rate constants and the corresponding pseudo-first-order rate constants at pH 8.0 for the reactions of phenolic compounds with the three oxidants. Results show that, even when bromide is present at very low concentrations compared with chlorine, rates of bromination of phenols are about 102–103fold greater than rates of chlorination. For phenol, corresponding half-life times are 10 min for chlorine and 300 ms for bromine. These results confirm that active bromine is probably rapidly scavenged by phenol-like organic structures incorporated into the natural organic matrix leading to the formation of brominated organic compounds during chlorination. For a concentration of ozone of 20 mM, pseudo-firstorder rate constants are 360 and 680 s1 for phenol and 4-chlorophenol, respectively (Table 2). Corresponding half-life times are about 2 and 1 ms, respectively. Comparison between ozone and bromine for both phenols shows that the pseudo-first-order rate constants for ozone are two to three orders of magnitude higher than the pseudo-first-order rate constants of bromine. During ozonation, the high reactivity of O3 with phenols limits the scavenging of active bromine by phenolic compounds. These results confirm previous findings which demonstrated that the formation of bromoorganic compounds is negligible during ozonation of natural waters [16].
5.5. Implications for drinking water treatment In order to estimate the behavior of HOBr in drinking water treatment plants, the pseudo-first-order rate constants of reaction between HOBr and phenol and 4-chlorophenol were calculated at pH 8.0 for a total concentration of HOBr of 1 mM (about 80 mg/L
6. Conclusion The objective of this study is the determination of rate constants of reactions between bromine and six phenolic
H. Gallard et al. / Water Research 37 (2003) 2883–2892
compounds in aqueous solution. *
*
*
*
The pH profiles of the apparent second-order rate constants can be explained by considering the individual reactions of HOBr with both undissociated and dissociated forms of phenol and the reaction of BrO with phenoxide ion. The reaction rate constants of HOBr with phenoxide ions are one to four orders of magnitude higher than the other individual reaction rate constants. This reaction controls the overall reaction rate between pH 6 and pH 10. The reaction rate constants of HOBr with phenoxide ions ranged between 1.4 103 and 2.1 108 M1 s1 for 2,4,6-trichlorophenol to pcresol, respectively. Hammett-type correlation was successfully tested for reaction of HOBr with phenoxide ions for phenolic compounds substituted in para and ortho positions. The Hammett equation (log k ¼ 8:023:33 (70.45)Sso;m;p ; r2 ¼ 0:946; n ¼ 5) can be used to estimate the reactivity of active bromine with phenolic compounds substituted in ortho or para positions. Comparison between the reactivity of bromine and oxidants used in drinking water treatment works showed that bromine is several orders more reactive than chlorine and iodine, but is less reactive than ozone with phenolic compounds. This confirmed that brominated disinfection by-products are rapidly produced during chlorination of NOM but that HOBr is probably not scavenged by organic compounds during ozonation in drinking water treatment plants.
[7] [8]
[9]
[10]
[11]
[12]
[13]
[14]
[15] [16]
References [1] Wajon JE, Morris JC. Bromination chemistry: rates of formation of NH2Br and some n-bromamino acids. In: Jolley et al., editors. Water chlorination: environmental impact and health effects, vol. 3. Ann Arbor, MI: Ann Arbor Science Publishers, Inc., 1980. pp. 171–181 [chapter 16]. [2] Trussell R, Umphres D. The formation of trihalomethanes. J Am Water Works Assoc 1978;11: 606–12. [3] Rook J, Gras A, van der Heijden G, de Wee J. Bromide oxidation and organic substitution in water treatment. J Environ Sci Health 1978;A13(2):91–116. [4] Boorman GA, Dellarco V, Dunnick JK, Chapin RE, Hunter S, Hauchman F, Gardner H, Cos M, Silss R. Drinking water disinfection byproducts: review and approach to toxicity evaluation. Environ Health Perspect 1999;107(suppl 1):207–17. [5] Morris RD. Drinking water and cancer. Environ. Health Perspect 1995;103(Suppl. 8):225–31. [6] Nieuwenhuijsen MJ, Toledano MB, Eaton NE, Fawell J, Elliott P. Chlorination disinfection by-products
[17]
[18]
[19]
[20]
[21]
[22]
2891
in water and their association with adverse reproductive outcomes: a review. Occup Environ Med 2000;57: 73–85. Legube B. A survey of bromate ion in European drinking water. Ozone Sci Eng 1996;18:325–48. von Gunten U, Oliveras Y. Advanced oxidation of bromide-containing waters: bromate formation mechanisms. Environ Sci Technol 1998;32:63–70. Westerhoff P, Song R, Amy G, Minear R. Numerical kinetic models for bromide oxidation to bromine and bromate. Water Res 1998;32(5):1687–99. Merlet N, de Laat J, Dor!e M. Oxydation des bromures au cours de la chloration des eaux de surface, incidence sur la production de compos!es organohalog!en!es. Revue Fr Sci Eau 1982;1:215–31. Symons M, Krasner W, Simms A, Sclimenti M. Measurement of THM and precursor concentrations revisited: the effect of bromide ion. J Am Water Works Assoc 1993;85:51–62. Boyce D, Barefoot C, Britton R, Horning F. Formation of trihalomethanes from the halogenation of 1,3-dihydroxybenzenes in dilute aqueous solution: synthesis of 2-13Cresorcinol and its reaction with chlorine and bromine. In: Water chlorination environmental and health effects, vol. 4. Ann Arbor, MI: Ann Arbor Science Publishers, 1983. p. 253–267. Bell RP, Rawlinson DJ. Kinetics of the bromination of some anisoles and phenols. J Chem Soc 1961: 63–8. Tee OS, Paventi M, Bennett JM. Kinetics and mechanism of the bromination of phenols and phenoxide ions in aqueous solution. Diffusion-controlled rates. J Am Chem Soc 1989;111:2233–40. de La Mare PBD. Electrophilic halogenation. London: Cambridge University Press, 1976. Pinkernell U, von Gunten U. Bromate minimization during ozonation: mechanistic considerations. Environ Sci Technol 2001;35(12):2525–31. Bischel Y, von Gunten U. Formation of iodotrihalomethanes during disinfection and oxidation of iodide-containing waters. Environ Sci Technol 2000;34: 2784–91. Gallard H, von Gunten U. Chlorination of phenols: kinetics and formation of chloroform. Environ Sci Technol 2002;36:884–90. Pinkernell U, Nowack B, Gallard H, von Gunten U. Methods for the photometric determination of reactive bromine and chlorine species with ABTS. Water Res 2000;34(18):4343–50. Troy RC, Margerum DW. Nonmetal redox kinetics— hypobromite and hypobromous acid reactions with iodide and with sulfite and the hydrolysis of bromosulfate. Inorg Chem 1991;30:3538–43. Palin AT. Current DPD methods for residual halogen compounds in water. J Am Water Works Assoc 1975; 67:32–3. Saguinsin JLS, Morris JC. The chemistry of aqueous nitrogen chloride. In: Johnson JD, editor. Disinfection water and wastewater. Ann Arbor, MI: Ann Arbor Science Publishers, Inc., 1975.
2892
H. Gallard et al. / Water Research 37 (2003) 2883–2892
[23] Rebenne LM, Gonzalez AC, Olson TM. Aqueous chlorination kinetics and mechanism of substituted dihydroxybenzenes. Environ Sci Technol 1996;30: 2235–42. [24] Hansch C, Leo A, Taft RW. A survey of Hammett substituent constants and resonance and field parameters. Chem Rev 1991;91:165–95.
[25] Hoign!e J, Bader H. Rate constants of reactions of ozone with organic and inorganic compounds in water. Water Res 1983;17:185–94. [26] Lee GF, Morris JC. Kinetics and chlorination of phenol— chlorophenolic tastes and odors. Int J Air Water Pollut 1962;6:19–431.
Rate constants of reactions of bromine with phenols in aqueous solution Herve! Gallard*, Fabien Pellizzari, Jean Philippe Croue! , B. Legube Laboratoire de Chimie de l’Eau et de l’Environnement, UMR CNRS 6008, Ecole Sup!erieure d’Ing!enieurs de Poitiers 40, Avenue du Recteur Pineau, Poitiers Cedex 86 022, France Received 15 May 2002; accepted 20 February 2003
Abstract The kinetics of bromination of six ortho- and para-substituted phenols was investigated between pH 5 and pH 12 in aqueous solution. Kinetics was followed with a continuous-flow reactor previously validated by studying the fast reaction between chlorine and ammonia. The overall reaction rate between bromine and phenols is controlled by the reaction of HOBr with the phenoxide ion between pH 6 and pH 10. The reaction of HOBr with the undissociated phenols and the reaction of BrO with the phenoxide ions become only significant for pHo6 and pH>10, respectively. The second-order rate constants for the reaction of HOBr with phenoxide ions vary between 1.4(70.1) 103 and 2.1(70.5) 108 M1 s1 for 2,4,6-trichlorophenol and 4-methylphenol, respectively. Hammett-type correlation was obtained for the reaction of HOBr with the phenoxide ions (log(k)=8.03.33 Ss) and was compared with Hammett-type correlations of HOCl and HOI. The reaction rate of bromine with phenol-like organic compounds was estimated to be about 103-fold higher than with chlorine and 103-fold lower than with ozone in drinking water treatment conditions. r 2003 Elsevier Science Ltd. All rights reserved. Keywords: Bromination; Phenols; Rate constants; Disinfection; Drinking water
1. Introduction In drinking water treatment plants, hypobromous acid (HOBr) is formed as a disinfectant by-product when chlorine or ozone is used to treat bromidecontaining natural waters. Once formed, hypobromous acid dissociates into hypobromite ion (BrO) with a pKa of 8.89: HOBr"BrO þ Hþ ;
Ka :
ð1Þ
*Corresponding author. Tel.: +33-5-49-45-44-31; fax: +335-49-45-97-68. E-mail address: [email protected] (H. Gallard).
In the presence of ammonia hypobromous acid reacts in a fast reaction producing monobromamine (NH2Br) and dibromamine (NHBr2) [1]. Hypobromous acid reacts also with natural organic matter (NOM) in complex reactions leading to the formation of bromoorganic compounds. The formation of brominated and mixed chlorobromo by-products such as bromoform (CHBr3) and mixed trihalomethanes (THMs) (dichlorobromomethane (CHCl2Br), dibromochloromethane (CHBr2Cl)) have been reported in finished drinking waters since the early 1970s [2,3]. Toxicological studies conducted on laboratory animals gave the first findings linking brominated by-products to colon cancer [4]. According to Morris [5], 5000 cases of bladder cancer per year and 8000 cases of rectal cancer per year may be associated with the
0043-1354/03/$ - see front matter r 2003 Elsevier Science Ltd. All rights reserved. doi:10.1016/S0043-1354(03)00132-5
2884
H. Gallard et al. / Water Research 37 (2003) 2883–2892
consumption of chlorinated drinking water in the US. The interest on trihalomethanes increased recently when epidemiological studies showed increased risks for birth defects, low birth weight, and miscarriages associated with the presence of THMs in drinking waters [6]. In the ozonation step, hypobromous acid reacts with ozone and OH radical. Hypobromous acid is the key intermediate in the formation mechanism of bromate (BrO 3 ), a potential carcinogenic by-product of ozonation in waterworks [7,8]. Kinetics of reaction of HOBr with ammonia, ozone, H2O2, and radicals has been extensively studied. The rate constants are now rather well known and the formation of bromate has been modeled in synthetic water [8]. In natural waters, the task appears more difficult because NOM can act as a scavenger of HOBr. To understand and evaluate bromate formation in natural waters, it was then concluded that it is necessary to better understand the NOM reactions with oxy-bromine species [9]. Number of studies has shown that the presence of bromide during chlorination increases the yield and the rate of formation of THMs [10,11,2]. To better understand the reactivity of chlorine or bromine with NOM, chlorination and, to a lesser extent, bromination of model organic compounds of NOM, such as phenolic compounds, have been studied. Bromination of resorcinol and orcinol yielded mono-, di- and tribrominated electrophilic substitution products in meta and para positions [12]. It is expected that 1,3-dihydroxy aromatic substrates would react with chlorine and bromine via the same pathway leading to the formation of chloroform and bromoform, respectively. The kinetics of bromination of phenols was studied by Bell and Rawlinson [13] and Tee et al. [14] at pHo7.0 in the presence of high concentration of bromide (2.5 mMo[KBr]o1.5 M). In such conditions, dibromine (Br2) and tribromide ion (Br 3 ) were considered as the only brominating species. Tee et al. [14] obtained, for phenol, a value of 1.2 109 M1 s1 for the attack of bromine on phenoxide ion and 8.5 108 M1 s1 for the reaction with tribromide ion. In the absence of bromide, the rate of bromination of aromatic compounds is proportional to the acidity of the media at low concentrations of acids [15]. It was hypothesized that the rate-limiting step is either the reaction of the positive bromine species H2OBr+ with organic compounds or the decomposition of a protonated transition state formed between HOBr and the organic substrate. An apparent second-order rate constant k ¼ 500 M1 s1 for the reaction between HOBr and phenol was determined in aqueous solution at pH 4.0 in batch reactor [16]. However, it is important to point out that for all these studies, pHs are too
low or the concentrations of bromide are too high to describe the reactivity of bromine with organic compounds in drinking water treatment. The kinetics of reaction of HOI and HOCl with phenols was studied in aqueous solutions between pH 2 and pH 12 [17,18]. The second-order rate constants of the reaction showed a pH-dependence that was explained by the speciation of both halogen and phenol species. For chlorine, apparent rate constants showed a maximum in the pH region 8–9, which is explained if HOCl is the only active electrophile and the reactivity of OCl is negligible. For hypobromous acid, a similar behavior is expected with a maximum shifted to higher pH because of a higher pKa for bromine. Few data are available regarding the kinetics of bromine in aqueous solution at near neutral pH in the absence of bromide. The purpose of this study was to determine the rate constants of initial bromination of selected phenolic compounds between pH 5 and pH 12.
2. Experimental section All chemicals were of the purest available quality. Solutions of organic compounds, ammonia, chlorine and bromine were prepared with ultra-pure water produced from a MilliQ (Millipore) water purification system. A total of six phenolic compounds was studied: phenol, p-cresol, p-chlorophenol, p-acetylphenol, 2,4dichlorophenol and 2,4,6-trichlorophenol. Stock solutions of chlorine were prepared by diluting a commercial solution of sodium hypochlorite (NaOCl, 4% active chlorine, Aldrich). Sodium hypochlorite was standardized by iodometry. Stock solutions of bromine were prepared from a solution of ozone (5 C, pH 4.0, 10 mM phosphate) by addition of 0.1 M solution of potassium bromide (KBr) [19]. After a day, ozone was purged by bubbling nitrogen for 30 min. The bromine solution was standardized by direct photometric determination of the hypobromite ion (BrO) at 329 nm (e ¼ 332 M1 cm1 [20]) after adjusting the pH of the solution to 11. For a ratio Br/O3=0.9, the yield of the reaction was X95% and the concentration of the bromine solution was typically 0.65 mM. The solution was stable for 4 weeks when stored at 4 C. All experiments were performed at room temperature (2372 C). pH measurements were carried out with a Tacussel PHM210 pH meter, which was calibrated with pH 4, 7 and 9 standard buffers (Merck). Solutions were buffered using phosphate (5 mM). Spectrophotometric measurements were performed on a SAFAS 190 DES spectrophotometer (SAFAS, Monaco).
H. Gallard et al. / Water Research 37 (2003) 2883–2892
A preliminary study showed that the initial reaction of bromine with most of the phenolic compounds used in this study was too fast to be followed in a batch reactor. Therefore, the kinetics of bromination was studied using both batch reactor and continuous-flow reactor.
3. Batch reactor experiments The batch reactor consisted of a 1-L glass bottle equipped with a dispenser. Experiments were initiated by adding an aliquot of a stock solution of 2,4,6trichlorophenol with micropipettes to the brominecontaining aqueous solution. At different reaction times, 3 mL of solution was rapidly transferred into a vial containing 0.1 mL of sodium sulfite to quench the reaction. Concentrations of 2,4,6-trichlorophenol were determined by HPLC (column ResolvTM 5 mm Spherical C18 waters; flow rate 1 mL/min; injection volume 200 mL; mobile phase 60% methanol/40% water/0.1% acetic acid; UV detection at 215 nm).
5.1. Validation of the continuous-flow apparatus The confirmation of the kinetics of the reaction between chlorine and ammonia was used to validate the measurements made using the continuous-flow apparatus. The reaction between hypochlorous acid (HOCl) and ammonia (NH3) is a fast, second-order reaction (reaction (2))—first-order with respect to chlorine and first-order with respect to ammonia. At 20 C, the reaction rate constant, kcl ; between the two neutral molecules is 5.6 106 M1 s1 [22]: HOCl þ NH3 -products; kcl :
For the other five selected phenolic compounds the experimental setup consisted of two HPLC pumps (Constametrics3200 LDC Analytical). The first one was used to pump the solution of oxidant (bromine or chlorine), the second one the solution of the tested molecule (phenolic compounds or ammonia). Both solutions were pumped with the same flow rate. Total flow rates in the reactor varied between 2 and 18 mL min1. The two solutions were mixed using a 1/ 1600 mixing tee, connected to a 78.0 cm or a 28.1 cm reaction tube (0.25-mm-ID Peek tube) where the reaction took place. The reaction solution was collected in a 5-cm long photometric cell containing 3 mL of a buffered (pH 6.4) solution of DPD. The exact collected volume was determined using a precision balance (70.1 mg). The DPD solution stopped the reaction and allowed analysis of the residual bromine (or chlorine) in solution at 510 nm [21]. Detection limit was 0.09 mM. The reaction time was defined based on the volume of the tube and the flux of the mixed solution. Reaction times ranged about from 10 to 150 ms.
5. Results and discussion A preliminary study was conducted to test the suitability of the continuous-flow apparatus for the study of fast kinetics.
ð2Þ
Both concentrations of HOCl and NH3 depend on pH according to equilibria (3) and (4): HOCl"ClO þ Hþ ; þ NHþ 4 "NH3 þ H ;
pKa2 ¼ 7:54; pKa3 ¼ 9:3:
ð3Þ ð4Þ
Considering [Cl2]t=[HOCl]+[OCl ] and [NH3]t= [NH3]+[NH+ 4 ] and both equilibria (3) and (4), the overall rate of chlorine consumption can be described as follows:
4. Continuous-flow apparatus experiments
2885
d½Cl2 t ¼ kapp ½Cl2 t ½NH3 t ; dt
ð5Þ
where kapp ¼ kcl ; ½Hþ
Ka3 : ½H þ Ka2 ½Hþ þ Ka3 þ
ð6Þ
The apparent reaction-rate constant, kapp ; based on the total concentration of chlorine and ammonia depends on the individual reaction-rate constant kcl ; pH and the equilibrium constants Ka2 and Ka3 (Eq. (6)). The experimental pH profile of kapp was determined between pH 5.6 and pH 10.6 under pseudo-first-order kinetic experimental conditions ½NH3 t b½Cl2 t : The initial concentration of ammonia varied between 0.1 and 2.4 mM. The initial concentration of chlorine was 5 mM. Inset in Fig. 1 shows, for three different experimental conditions, the plot of Ln ½Cl2 t =½Cl2 t0 vs. time (to150 ms). The observed pseudo-first-order rate constant, kobs ; was determined as the slope of the linear curve: ½Cl2 t Ln ¼ kobs t; ð7Þ ½Cl2 t0 where the observed rate constant kobs ¼ kapp ½NH3 t0 :
ð8Þ
Fig. 1 depicts the experimental and theoretical profile of the apparent second-order rate constant (kapp ) vs. pH. The theoretical pH profile was determined by a non-linear least-squares regression of experimental data using Eq. (6) and the fit curve function of SigmaPlot 2000 software as a tool. The calculated value of kcl was 5.97(70.54) 106 M1 s1. This value ob-
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2886
4 Ln ([Cl2]t /[Cl2]t0)
-1
Log (kapp) (M s-1)
5
3
0 -0.2 -0.4 -0.6 -0.8 -1 -1.2
pH = 7.1 pH = 7.4
0
0.05
pH = 8.9 0.1
0.15
Time (sec.)
2 5
6
7
8 pH
9
10
11
Fig. 1. Apparent rate constant pH profile and pseudo-firstorder kinetic plot (inset) of the reaction of chlorine with ammonia. Conditions: 23(72) C, 0.1 mMo[NH3]t0o2.4 mM, [Cl2]t0=5 mM.
tained with the continuous-flow apparatus is close to the value of 5.6 106 M1 s1 given by Saguinsin and Morris [22]. It was concluded that this experimental setup can be conveniently used for studying the fast reactions of halogens, e.g. the reaction of bromine with phenols. 5.2. Determination of reaction rate constants between bromine and highly reactive phenolic species Based on the kinetic study of iodination [17] and chlorination [23,18] at neutral and alkaline pH, the initial bromination of phenolic compounds was considered as a second-order reaction (reaction (9))—firstorder in active bromine ([Br2]t=[HOBr]+[BrO]) and first-order in phenol ([Ph]t=[PhOH]+[PhO]): Ph þ Br2 -Products; k:
ð9Þ
Considering the speciation of both bromine (reaction (1), pKa ¼ 8:89) and phenolic compound (reaction (10)), the second-order rate constant, k; based on the total concentration of active bromine and phenolic compound, will depend on the degree of dissociation of all species: PhOH$PhO þ Hþ ;
Ka4 :
ð10Þ
Between pH 5 and pH 12, hypobromous acid is probably the most reactive species and can react with both undissociated and dissociated phenolic species according to reactions (11) and (12), respectively. The hypobromite ion is less reactive and might only react with phenoxide ion in alkaline solutions (reaction (13)): HOBr þ PhOH-Products; k1 ;
ð11Þ
HOBr þ PhO -Products;
k2 ;
ð12Þ
BrO þ PhO -Products;
k3 :
ð13Þ
In the first assumption, the dibromine (Br2) species was not introduced into the model. Considering the yield of 95% for the conversion of Br into active bromine with ozone (see experimental section) and an equilibrium constant of 5.8 109 for reaction (14), the proportion of dibromine was calculated to be o0:08% at pH 5. Its reaction with phenol was then considered negligible because the kinetic constants of HOBr reported later in the paper are only one order of magnitude smaller than the Br2 rate constant determined by Tee et al. [14] for the same compound; results that reinforce our hypothesis: HOBr þ Br "Br2 þ OH :
ð14Þ
The reaction between OBr and PhOH was considered to be negligible, because when introduced into the model, the best fit was obtained when the k value tends to zero. Also, one should note that in the literature the reactions between OX and phenol species are generally not included into models. Our results showed that the contribution of the OBr/PhO reaction seems to be significant only above pH 11. Consequently, the consumption of active bromine can be written as follows: d½Br2 t ¼ k1 ½HOBr ½PhOH k2 ½HOBr ½PhO
dt k3 ½BrO ½PhO ;
ð15Þ
d½Br2 t ¼ ððk1 ð1 aBr Þð1 aPh Þ þ k2 ð1 aBr ÞaPh dt þ k3 aBr aPh ÞÞ½Br2 t ½Ph t ; ð16Þ kapp ¼ k1 ð1 aBr Þð1 aPh Þ þ k2 ð1 aBr ÞaPh þ k3 aBr aPh ; ð17Þ where aBr and aPh are the degrees of dissociation of hypobromous acid (aBr =[BrO]/([HOBr]+[BrO]) and phenolic compound (aPh =[PhO]/([PhO]+[PhOH])). The degrees of dissociation can be calculated from the pH and the dissociation constants as follows: aBr ¼
Ka1 ; ðKa1 þ ½Hþ Þ
ð18Þ
aPh ¼
Ka4 : ðKa4 þ ½Hþ Þ
ð19Þ
Replacing aBr and aPh by Eqs. (18) and (19) into Eq. (17), the second-order rate constant, kapp ; depend on the individual reaction-rate constants (k1 ; k2 ; k3 ), the pH, and the acid dissociation constants according to
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2887
Table 1 pKa ; Hammett constants and rate constants of HOBr (k1 ; k2 ) and OBr (k3 ) with selected phenolic compounds Phenolic compounds
pKa
Ss
k1 (M1 s1)
k2 (M1 s1)
k3 (M1 s1)
p-Cresol Phenol 4-Chlorophenol 4-Acetylphenol 2,4-Dichlorophenol 2,4,6-Trichlorophenol
10.26 10.0 9.43 8.6 7.85 6.17
0.13 0 0.24 0.47 0.60 0.97
o1.0 104 o500 6.0(73.0) 103 1.0(70.4) 104 3.0(71.0) 104 —
2.1(70.5) 108 1.8(70.2) 108 7.0(70.8) 106 4.1(70.6) 106 8.8(70.9) 105 1.4(70.1) 103
o3.0 105 o1.0 105 5.5(74.0) 104 o2.0 103 3.0(72.0) 104 p1
0
6 pH 12.0 -1 -1
Log (kapp) en M s
-1 pH 8.9
-1.5
pH 10.1
-2
150
kobs (s -1)
Ln ([Br2]t /[Br2]to)
-0.5
-2.5
100 50
5
0
-3
0
20
40
60
[dichlorophenol]to (µM)
4.5
-3.5 0
0.05
0.1
0.15
Time (sec.)
Fig. 2. Pseudo-first-order kinetic plots of the bromination of 2,4-dichlorophenol. Conditions: 23(72) C, [2,4-dichlorophenol]t0=125 mM, [Br2]t0=8 mM. The inset shows the evolution of kobs with the initial concentration of 2,4-dichlorophenol at pH 7.9.
Eq. (20): þ 2
kapp ¼
5.5
þ
k1 ½H þ k2 ½H Ka4 þ k3 Ka1 Ka4 : ðKa1 þ ½Hþ ÞðKa4 þ ½Hþ Þ
ð20Þ
The individual rate constants were determined by a nonlinear least-squares regression of the experimental pH profile of kapp : The kinetics of initial bromination of phenol, 4chloro, 2,4-dichlorophenol, 4-acetylphenol and p-cresol were studied between pH 5 and pH 12 with the continuous-flow apparatus under experimental conditions developed for pseudo-first-order kinetics ð½Br2 t0 515 ½Ph t0 Þ: The acid dissociation constants of the studied phenols are given in Table 1. The initial concentrations of phenols varied from 25 to 125 mM. The initial concentrations of bromine ranged from 1.3 to 7.9 mM. In the present paper, the bromination of 2,4-dichlorophenol is described in detail and the results for the other selected phenols are presented less extensively.
4
5
6
7
8
9
10
11
12
13
pH Fig. 3. pH profile of the second-order apparent rate constant of initial bromination of 2,4-dichlorophenol. Conditions: 23(72) C, 25 mMo[2,4-dichlorophenol]t0o125 mM, 1.3 mMo [Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
Fig. 2 plots for three different pH values the evolution of Ln ½Br2 t =½Br2 t0 vs. time for the bromination of a 125 mM solution of 2,4-dichlorophenol. Results show that the rate of bromination can be well linearized using a pseudo-first-order kinetic model. The reaction order with respect to 2,4dichlorophenol was examined at pH 7.9 by varying the concentration of 2,4-dichlorophenol. A linear plot could be drawn between the pseudo-first-order rate constant and the initial concentration of 2,4dichlorophenol (inset in Fig. 2). Therefore, the rate of the bromine disappearance is first order in the concentrations of bromine and 2,4-dichlorophenol. Fig. 3 shows the experimental and theoretical pH profiles of kapp for the bromination of 2,4-dichlorophenol between pH 5 and pH 12. As for chlorine [18,23] and iodine [17], pH profiles exhibit a maximum in near-alkaline media, which is explained if HOBr is the major electrophilic species. Using the acid dissociation
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2888 7
-1
-1
Log (kapp) (M s )
6.5 6 p -methylphenol
5.5 5
p -chlorophenol
4.5 4 3.5 3 4
5
6
7
8 9 pH
10
11
12
13
Fig. 4. pH profile of the second-order apparent rate constant of initial bromination of p-methylphenol and p-chlorophenol. Conditions: 23(72) C, 25 mMo[Ph]t0o125 mM, 1.3 mMo [Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
7 6.5
phenol
-1
-1
Log (kapp) (M s )
6 5.5 5 4.5 p -acetylphenol
4 3.5
the pH profile of kapp obtained in our study for phenol is in good agreement with the apparent second-order rate constant found by Pinkernell and von Gunten [16] at pH 4.0 in a batch reactor (kapp ¼ 500 M1 s1). The individual reaction-rate constants of phenol and substituted phenols studied with the continuous-flow apparatus are listed in Table 1. The rate constants of HOBr with the phenoxide ions range from 2.1 108 M1 s1 for p-cresol to 8.8 105 M1 s1 for 2,4-dichlorophenol. For phenol, the rate constant of HOBr with phenoxide ion is 1.8 108 M1 s1, which is about one order of magnitude smaller than the corresponding rate constant for the Br2 species (k ¼ 1:2 109 M1 s1 [14]). This verified that, in our condition, the rate constant with dibromine is not enough to compensate the low Br2 concentrations. In agreement with the electrophilic character of HOBr the rate constants of HOBr with the phenoxide ions decrease with the deactivating effect of the substituents. These rate constants are one to four orders of magnitude greater than the corresponding rate constants of the reaction between HOBr and the undissociated phenols and the rate constants of the reaction between BrO with the phenoxide ions. In pH region 6–9, the reaction between HOBr and the phenoxide ions controls the overall reaction. As a consequence, the precision for k1 and k3 is very low and only upper limits were determined for phenol, pcresol and p-acetylphenol. 5.3. Bromination of 2,4,6-trichlorophenol
3 2.5 2 4
5
6
7
8
9
10
11
12
13
pH
Fig. 5. pH profile of the second-order apparent rate constant of initial bromination of phenol and p-acetylphenol (full and open circles: this study; open square: from [16]). Conditions: 23(72) C, 25 mMo[Ph]t0o125 mM, 1.3 mMo[Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k1 ; k2 and k3 to the experimental data (symbols).
constants Ka1 ¼ 108:9 and Ka4 ¼ 107:85 the individual rate constants were calculated as k1 ¼ 3:0ð71:0Þ 104 M1 s1, k2 ¼ 8:8ð70:9Þ 105 M1 s1 and k3 ¼ 3:0ð72:0Þ 104 M1 s1. The same approach was used to model the bromination kinetics of the other selected phenolic compounds. Figs. 4 and 5 show the pH dependency of the secondorder rate constant for both 4-methylphenol and 4chlorophenol (Fig. 4) and for both phenol and 4acetylphenol (Fig. 5). Again, the proposed model fits well with the experimental profile. Fig. 5 shows also that
Since 2,4,6-trichlorophenol exhibits a much lower reactivity toward active bromine, kinetics of bromination were studied in a classical batch reactor under pseudo-first-order experimental conditions ([2,4,6-trichlorophenol]t05[Br2]t0). The initial concentration of 2,4,6-trichlorophenol was 0.5 or 1 mM. The concentration of active bromine was >20 mM. Experiments were conducted between pH 5.0 and pH 11.2. Between pH 7.0 and pH 11.2 the pseudo-firstorder rate constant determined from the linear plots ln[trichlorophenol]t/[trichlorophenol]t0 vs. time was found to be directly proportional to the initial concentration of 2,4,6-trichlorophenol (results not presented). The disappearance of 2,4,6-trichlorophenol is first order with respect to bromine and 2,4,6-trichlorophenol concentrations. A rate constant of 1400(7100) M1 s1 was found for the reaction of HOBr with the 2,4,6trichlorophenoxide ion by fitting experimental values with the above proposed model. Both experimental and theoretical pH profiles are plotted in Fig. 6. In this pH region, the overall reaction is only controlled by the reaction of HOBr with the 2,4,6-trichlorophenoxide ion. The rate constant of the reaction between BrO and 2,4,6-trichlorophenolate was estimated to be below
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2889
10 3
Log (k2) (M s-1)
2 -1
-1
-1
Log (kapp) (M s )
9
1 0
p- Me
Ph
8 7
HOI
y = -3.33 (±0.45) . Σσ + 7.98 2 R = 0.95
p- Cl
p -Ac
6
2,4-diCl
5
HOCl
4
-1
2,4,6-triCl
3 -2 4
6
8 pH
10
12
Fig. 6. pH profile of the second-order apparent rate constant of initial bromination (( ) this study) and chlorination [26] of 2,4,6-trichlorophenol. Conditions: 23(72) C, 25 mMo [Ph]t0o125 mM, 1.3 mMo[Br2]t0o8 mM. The solid line is a non-linear least-squares regression fit of k2 and k3 to the experimental data (symbols).
1 M1 s1. For comparison the pH profile of the secondorder rate constant of chlorination of 2,4,6-trichlorophenol is also plotted in Fig. 6. Results show that the rate constants of bromination are more than two orders of magnitude greater than the rate constants of chlorination. As expected, the maximum of the pH profile is shifted towards higher pH values from chlorine to bromine due to the higher pKa of HOBr/BrO. For pHo7 and in our experimental conditions, the reaction rate between bromine and 2,4,6-trichlorophenol deviated from a second-order kinetic model. At pH 5.8, the degradation rate of 2,4,6-trichlorophenol could be linearized using a pseudo-first-order kinetic model; however, the degradation rate did not depend on the concentration of bromine. The rate of trichlorophenol disappearance was found to be zero order with respect to [HOBr]. In such conditions, the reaction rate of HOBr with the trichlorophenoxide ion is not the ratelimiting step. The rate-limiting step is the deprotonation of 2,4,6-trichlorophenol and the observed pseudo-first0 order rate constant, kobs ; is equal to the rate constant of deprotonation kd : A first-order rate constant 0 kd ¼ kobs ¼ 2:3 103 s1 was found. The rate constant of the reaction between the trichlorophenoxide ion and H+ can then be determined from the acid dissociation constant Ka4 : kf ¼ kd =Ka4 : With pKa4 ¼ 6:17 (Ka4 ¼ 6:8 107 ), a value kf ¼ 3400 M1 s1 was calculated for the reaction between 2,4,6-trichlorophenol and H+. 5.4. Linear free-energy relationships The Hammett correlation is classically used to predict the effect of substituents on reaction-rate constants
2 -0.5
0
0.5 Σσ
1
1.5
Fig. 7. Hammett plots for the reaction of HOBr, HOI and HOCl with phenoxide ions (HOBr: this study; HOCl: from [18] HOI: from [17]). p-Me: p-methylphenol; Ph: phenol; p-Cl: pchlorophenol; p-Ac: p-acetylphenol; 2,4-diCl: 2,4-dichlorophenol; and 2,4,6-triCl: 2,4,6-trichlorophenol.
between organic compounds and oxidants [23,17,18]. The Hammett constant (s) reflects the effects of substituents on the electron density of the aromatic ring by inductive and resonance effects. For s > 0; substituents have an electron-withdrawing effect (e.g. CN), whereas for so0 substituents exert an electron-donating effect (e.g. CH3). The non-substituted phenol was used as a reference compound (Sso;m;p ¼ 0). Values of s were taken from Hansch et al. [24] and the values of Sso;m;p are shown in Table 1 for the studied phenolic compounds. The Hammett correlation was only tested for the reaction between HOBr and phenoxide ions because these rate constants were determined with the best precision and also because this particular reaction controls the reactivity of bromine in natural waters (neutral pH). Fig. 7 shows that a good linear plot is obtained for log(k2 ) vs. Sso;m;p with our results. The linear regression for the reaction of bromine with the five selected phenolic compounds substituted in para or/ and in ortho position is X logðk2 Þ ¼ 8:0 3:33ð70:45Þ so;m;p ; r ¼ 0:946;
n ¼ 5:
ð21Þ
The negative sign of the Hammett slope (r) is typical of electrophilic substitution reaction because the increase of the Hammett constant corresponds to an increase of the deactivating effect of the substituents. With phenolic compounds the OH and O substituents act as activating substituents that promote the halogen substitution onto the ortho and para positions. The initial reaction of bromine
H. Gallard et al. / Water Research 37 (2003) 2883–2892
2890
Table 2 Rate constants and pseudo-first-order reaction rates for the reaction of HOBr, HOCl and O3 with phenol and 4-chlorophenol at pH 8.0 (2372 C) Phenolic compound
Phenol 4-Chlorophenol
HOCl
HOBr
O3
k (M1 s1)b
k0 (s1)a
k (M1 s1)c
k0 (s1)a
k (M1 s1)d
k0 (s1)a
60 30
1.2 103 6.0 104
2.2 106 2.3 105
2.2 0.23
18 106 34 106
360 680
a
Pseudo-first-order rate constants are calculated from second-order rate constants for [HOCl]=20 mM, [HOBr]=1 mM and [O3]=20 mM. b From [18]. c Our study. d From [25].
with the studied phenols is therefore the electrophilic substitution onto available 2, 4 and 6 ring positions leading to the formation of bromophenols [14]. As a consequence the above correlation cannot be used to predict the rate constant of bromine with 2,4,6trichlorophenoxide ion since, in this case, all para and ortho positions are already substituted with chlorine atoms. This is confirmed in Fig. 7 in which rate constant of 2,4,6-trichlorophenoxide ion is significantly below the Hammett correlation. Fig. 7 gives for comparison the Hammett-type correlations established for the rate constants of HOCl [18] and HOI [17] with phenoxide ions. Results show that HOBr is the most reactive oxy-halides. Rate constants of HOBr are about three to four orders of magnitude higher than for HOCl and between 5 and 500-fold higher than for HOI. The intermediate reactivity of iodine is not consistent with the order of electronegativity of halogens. Because iodine has the lowest electronegativity, hypoiodous acid should present the highest rate constants. These results remain unclear. One explanation may be the low stability of HOI in aqueous solution, which makes the determination of its rate constants with organic compounds more difficult. Steric effect might also explain the unexpected low reactivity of iodine. The Hammett slopes of HOCl (r ¼ 3:0) and HOBr (r ¼ 3:33) are rather similar but the rate constants of HOI show a stronger effect of substituents with r ¼ 5:4: As it can be expected from electronegativity scale, the jrj value increases with decreasing electronegativity of halogens.
of Br). Corresponding pseudo-first-order rate constants were calculated at pH 8.0 for ozone ([O3]=20 mM; D1 mg/L) and chlorine ([HOCl]=20 mM; D1.4 mg/L) from literature data. Table 2 presents the apparent second-order rate constants and the corresponding pseudo-first-order rate constants at pH 8.0 for the reactions of phenolic compounds with the three oxidants. Results show that, even when bromide is present at very low concentrations compared with chlorine, rates of bromination of phenols are about 102–103fold greater than rates of chlorination. For phenol, corresponding half-life times are 10 min for chlorine and 300 ms for bromine. These results confirm that active bromine is probably rapidly scavenged by phenol-like organic structures incorporated into the natural organic matrix leading to the formation of brominated organic compounds during chlorination. For a concentration of ozone of 20 mM, pseudo-firstorder rate constants are 360 and 680 s1 for phenol and 4-chlorophenol, respectively (Table 2). Corresponding half-life times are about 2 and 1 ms, respectively. Comparison between ozone and bromine for both phenols shows that the pseudo-first-order rate constants for ozone are two to three orders of magnitude higher than the pseudo-first-order rate constants of bromine. During ozonation, the high reactivity of O3 with phenols limits the scavenging of active bromine by phenolic compounds. These results confirm previous findings which demonstrated that the formation of bromoorganic compounds is negligible during ozonation of natural waters [16].
5.5. Implications for drinking water treatment In order to estimate the behavior of HOBr in drinking water treatment plants, the pseudo-first-order rate constants of reaction between HOBr and phenol and 4-chlorophenol were calculated at pH 8.0 for a total concentration of HOBr of 1 mM (about 80 mg/L
6. Conclusion The objective of this study is the determination of rate constants of reactions between bromine and six phenolic
H. Gallard et al. / Water Research 37 (2003) 2883–2892
compounds in aqueous solution. *
*
*
*
The pH profiles of the apparent second-order rate constants can be explained by considering the individual reactions of HOBr with both undissociated and dissociated forms of phenol and the reaction of BrO with phenoxide ion. The reaction rate constants of HOBr with phenoxide ions are one to four orders of magnitude higher than the other individual reaction rate constants. This reaction controls the overall reaction rate between pH 6 and pH 10. The reaction rate constants of HOBr with phenoxide ions ranged between 1.4 103 and 2.1 108 M1 s1 for 2,4,6-trichlorophenol to pcresol, respectively. Hammett-type correlation was successfully tested for reaction of HOBr with phenoxide ions for phenolic compounds substituted in para and ortho positions. The Hammett equation (log k ¼ 8:023:33 (70.45)Sso;m;p ; r2 ¼ 0:946; n ¼ 5) can be used to estimate the reactivity of active bromine with phenolic compounds substituted in ortho or para positions. Comparison between the reactivity of bromine and oxidants used in drinking water treatment works showed that bromine is several orders more reactive than chlorine and iodine, but is less reactive than ozone with phenolic compounds. This confirmed that brominated disinfection by-products are rapidly produced during chlorination of NOM but that HOBr is probably not scavenged by organic compounds during ozonation in drinking water treatment plants.
[7] [8]
[9]
[10]
[11]
[12]
[13]
[14]
[15] [16]
References [1] Wajon JE, Morris JC. Bromination chemistry: rates of formation of NH2Br and some n-bromamino acids. In: Jolley et al., editors. Water chlorination: environmental impact and health effects, vol. 3. Ann Arbor, MI: Ann Arbor Science Publishers, Inc., 1980. pp. 171–181 [chapter 16]. [2] Trussell R, Umphres D. The formation of trihalomethanes. J Am Water Works Assoc 1978;11: 606–12. [3] Rook J, Gras A, van der Heijden G, de Wee J. Bromide oxidation and organic substitution in water treatment. J Environ Sci Health 1978;A13(2):91–116. [4] Boorman GA, Dellarco V, Dunnick JK, Chapin RE, Hunter S, Hauchman F, Gardner H, Cos M, Silss R. Drinking water disinfection byproducts: review and approach to toxicity evaluation. Environ Health Perspect 1999;107(suppl 1):207–17. [5] Morris RD. Drinking water and cancer. Environ. Health Perspect 1995;103(Suppl. 8):225–31. [6] Nieuwenhuijsen MJ, Toledano MB, Eaton NE, Fawell J, Elliott P. Chlorination disinfection by-products
[17]
[18]
[19]
[20]
[21]
[22]
2891
in water and their association with adverse reproductive outcomes: a review. Occup Environ Med 2000;57: 73–85. Legube B. A survey of bromate ion in European drinking water. Ozone Sci Eng 1996;18:325–48. von Gunten U, Oliveras Y. Advanced oxidation of bromide-containing waters: bromate formation mechanisms. Environ Sci Technol 1998;32:63–70. Westerhoff P, Song R, Amy G, Minear R. Numerical kinetic models for bromide oxidation to bromine and bromate. Water Res 1998;32(5):1687–99. Merlet N, de Laat J, Dor!e M. Oxydation des bromures au cours de la chloration des eaux de surface, incidence sur la production de compos!es organohalog!en!es. Revue Fr Sci Eau 1982;1:215–31. Symons M, Krasner W, Simms A, Sclimenti M. Measurement of THM and precursor concentrations revisited: the effect of bromide ion. J Am Water Works Assoc 1993;85:51–62. Boyce D, Barefoot C, Britton R, Horning F. Formation of trihalomethanes from the halogenation of 1,3-dihydroxybenzenes in dilute aqueous solution: synthesis of 2-13Cresorcinol and its reaction with chlorine and bromine. In: Water chlorination environmental and health effects, vol. 4. Ann Arbor, MI: Ann Arbor Science Publishers, 1983. p. 253–267. Bell RP, Rawlinson DJ. Kinetics of the bromination of some anisoles and phenols. J Chem Soc 1961: 63–8. Tee OS, Paventi M, Bennett JM. Kinetics and mechanism of the bromination of phenols and phenoxide ions in aqueous solution. Diffusion-controlled rates. J Am Chem Soc 1989;111:2233–40. de La Mare PBD. Electrophilic halogenation. London: Cambridge University Press, 1976. Pinkernell U, von Gunten U. Bromate minimization during ozonation: mechanistic considerations. Environ Sci Technol 2001;35(12):2525–31. Bischel Y, von Gunten U. Formation of iodotrihalomethanes during disinfection and oxidation of iodide-containing waters. Environ Sci Technol 2000;34: 2784–91. Gallard H, von Gunten U. Chlorination of phenols: kinetics and formation of chloroform. Environ Sci Technol 2002;36:884–90. Pinkernell U, Nowack B, Gallard H, von Gunten U. Methods for the photometric determination of reactive bromine and chlorine species with ABTS. Water Res 2000;34(18):4343–50. Troy RC, Margerum DW. Nonmetal redox kinetics— hypobromite and hypobromous acid reactions with iodide and with sulfite and the hydrolysis of bromosulfate. Inorg Chem 1991;30:3538–43. Palin AT. Current DPD methods for residual halogen compounds in water. J Am Water Works Assoc 1975; 67:32–3. Saguinsin JLS, Morris JC. The chemistry of aqueous nitrogen chloride. In: Johnson JD, editor. Disinfection water and wastewater. Ann Arbor, MI: Ann Arbor Science Publishers, Inc., 1975.
2892
H. Gallard et al. / Water Research 37 (2003) 2883–2892
[23] Rebenne LM, Gonzalez AC, Olson TM. Aqueous chlorination kinetics and mechanism of substituted dihydroxybenzenes. Environ Sci Technol 1996;30: 2235–42. [24] Hansch C, Leo A, Taft RW. A survey of Hammett substituent constants and resonance and field parameters. Chem Rev 1991;91:165–95.
[25] Hoign!e J, Bader H. Rate constants of reactions of ozone with organic and inorganic compounds in water. Water Res 1983;17:185–94. [26] Lee GF, Morris JC. Kinetics and chlorination of phenol— chlorophenolic tastes and odors. Int J Air Water Pollut 1962;6:19–431.