Oxidative treatment of bromide containing waters formation of bromine and its reactios with inorganic and organic.pdf

8/19/2019 Oxidative treatment of bromide containing waters formation of bromine and its reactios with inorganic and organi…

1/28

Oxidative treatment of bromide-containing waters:Formation of bromine and its reactions withinorganic and organic compounds d A criticalreview

Michèle B. Heeba , Justine Criquet b ,c,d , Saskia G. Zimmermann-Steffens a ,Urs von Gunten a ,b ,e,*a School of Architecture, Civil and Environmental Engineering (ENAC), Ecole Polytechnique Fé dé rale de Lausanne(EPFL), CH-1015 Lausanne, Switzerlandb Eawag, Swiss Federal Institute of Aquatic Science and Technology, CH-8600 Du¨ bendorf, Switzerlandc Université Lille 1, Laboratoire Gé osystèmes, UMR CNRS 8217, 59655 Villeneuve d’Ascq, Franced Curtin Water Quality Research Centre, Curtin University, GPO Box U1987, Perth, WA 6845, Australiae Institute of Biogeochemistry and Pollutant Dynamics, ETH Zu¨ rich, CH-8092 Zü rich, Switzerland

a r t i c l e i n f o

Article history:Received 25 June 2013Received in revised form23 August 2013Accepted 25 August 2013Available online 4 September 2013

Keywords:BromineHOBrOxidation kineticsWater treatmentInorganic compoundsOrganic compounds

a b s t r a c t

Bromide (Br ) is present in all water sources at concentrations ranging from w 10 to> 1000 mg L 1 in fresh waters and about 67 mg L 1 in seawater. During oxidative watertreatment bromide is oxidized to hypobromous acid/hypobromite (HOBr/OBr ) and otherbromine species. A systematic and critical literature review has been conducted on thereactivity of HOBr/OBr and other bromine species with inorganic and organic compounds,including micropollutants.

The speciation of bromine in the absence and presence of chloride and chlorine hasbeen calculated and it could be shown that HOBr/OBr are the dominant species in freshwaters. In ocean waters, other bromine species such as Br 2, BrCl, and Br2O gain importanceand may have to be considered under certain conditions.

HOBr reacts fast with many inorganic compounds such as ammonia, iodide, sulte,nitrite, cyanide and thiocyanide with apparent second-order rate constants in the order of 104e 109 M 1 s 1 at pH 7. No rate constants for the reactions with Fe(II) and As(III) areavailable. Mn(II) oxidation by bromine is controlled by a Mn(III,IV) oxide-catalyzed processinvolving Br 2O and BrCl.

Bromine shows a very high reactivity toward phenolic groups (apparent second-orderrate constants kapp z 103e 105 M 1 s 1 at pH 7), amines and sulfamides ( kapp z 105e 106 M 1 s 1 at pH 7) and S-containing compounds ( kapp z 105e 107 M 1 s 1 at pH 7). Forphenolic moieties, it is possible to derive second-order rate constants with a Hammett- s -based QSAR approach with log ðkðHOBr=PhO ÞÞ ¼ 7:8 3:5Ss . A negative slope is typical forelectrophilic substitution reactions.

In general, kapp of bromine reactions at pH 7 are up to three orders of magnitude greaterthan for chlorine. In the case of amines, these rate constants are even higher than for

* Corresponding author . Eawag, Swiss Federal Institute of Aquatic Science and Technology, CH-8600 Du ¨ bendorf, Switzerland. Tel.: þ 41 58765 5270; fax: þ 41 58 765 5210.

E-mail addresses: [email protected] , [email protected] (U. von Gunten).

Available online at www.sciencedirect.com

journal homepage: www.elsevier.com/locate /watres

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2

0043-1354/$ e see front matter ª 2013 Elsevier Ltd. All rights reserved.http://dx.doi.org/10.1016/j.watres.2013.08.030

mailto:[email protected]:[email protected]://www.sciencedirect.com/science/journal/00431354http://www.elsevier.com/locate/watreshttp://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://www.elsevier.com/locate/watreshttp://www.sciencedirect.com/science/journal/00431354http://crossmark.crossref.org/dialog/?doi=10.1016/j.watres.2013.08.030&domain=pdfmailto:[email protected]:[email protected]


2/28

ozone. Model calculations show that depending on the bromide concentration and the pH,the high reactivity of bromine may outweigh the reactions of chlorine during chlorinationof bromide-containing waters.

ª 2013 Elsevier Ltd. All rights reserved.

1. Introduction

Oxidative water treatment such as chlorination or ozonationhas been applied in the rst half of the 20th century primarilyfor disinfection purposes. Consecutively, other treatmentgoals such as the removal of taste and odor compounds andthe elimination of (anthropogenic) micropollutants weretackled as well ( Sedlak and von Gunten, 2011 ). One majordrawback of oxidative water treatment is the formation of disinfection by-products (DBPs) and transformation productsresulting from the reaction of oxidants with moieties of thewater matrix and micropollutants, respectively, that might beof human or environmental health concern ( Krasner et al.,2006; von Sonntag and von Gunten, 2012 ).

Besides natural organic matter (NOM), which reacts withalloxidantsapplied in water treatment,bromide (Br )isoneof the key components of the water matrix relevant in oxidationprocesses; it can be oxidized to hypobromous acid (HOBr, Eqs.(1) and (7)) and other bromine species. Various studies haveillustrated that the presence of bromide during oxidationprocesses can have a signicant impact: the formation of bromine can accelerate the transformation of undesiredcompounds ( Haag et al., 1984; Lee and von Gunten, 2009 ),prevent the formationof undesiredby-products ( Criquet et al.,2012) and affect the extent of formation of halogenated DBPs(e.g., Amy et al., 1984; Bulloch et al., 2012; Chowdhury et al.,2010; Cowman and Singer, 1996; Gallard et al., 2003; Huet al., 2006; Hua et al., 2006; Inaba et al., 2006; Jones et al.,2012; Pan and Zhang, 2013; Richardson et al., 2003; Rodilet al., 2012; Rook et al., 1978; Singer, 1994; Symons et al.,1993; Zhang et al., 2005; Zhao et al., 2012 ) and N-nitroso-dimethylamine (NDMA) ( Le Roux et al., 2012; Luh and Marin ˜ as,2012; von Gunten et al., 2010 ). Bromide often acts as a catalystin these reactions ( Criquet et al., 2012; Duirk et al., 2008; Haag et al., 1984; von Gunten et al., 2010 ). In Fig. 1, a schematicoverview of the role of bromide in oxidative water treatmentis given.

This review provides an assessment of the role of bromine(the term bromine is used for all active bromine species,mainly HOBr, OBr , and Br2, but also BrCl, Br2O, and BrOCl)during oxidation processes with an emphasis on kinetics andmechanisms of bromine reactions. This information may alsohelp to better understand the formation of (brominated)disinfection by-products such as trihalomethanes (THMs) andhaloacetic acids (HAAs). However, their formation is not themain focus of this review. Based on literature data, an over-view of bromine speciation and bromine reactivity withinorganic and organic compounds is presented. For relevantinorganic substances (NH 3, NO2 , SO3 2 , I , CN , SCN ) andtypical functional groups, such as activated aromatic systems,amines, sulfur groups, and olenes, bromination kinetics andmechanisms are described. Furthermore, for some organic

micropollutants relevant for urban water management, adiscussion on expected and observed bromine reactivities isprovided. Finally, a kinetic assessment is made on thecontribution of bromine to the overall reaction during chlori-nation of bromide-containing waters.

2. Aqueous bromine chemistry

2.1. Bromide levels in natural waters

Bromide is present in all fresh waters in concentrations in therange of a few mg L 1 to several mgL 1 (Magazinovic et al.,2004; von Gunten and Hoigne ´, 1992) and is non-toxic atthese concentrations ( Flury and Papritz, 1993 ). The concen-tration of bromide in seawater is about 67 mg L 1 (Millero,1974). Local geological situations and salt-water intrusion incoastal areas are responsible for the observed bromide levelsin natural waters ( Agus et al., 2009). However, also anthropo-genic activities such as potassium and coal mining or chem-ical industries may result in elevated bromide levels ( Valeroand Arbó s, 2010; von Gunten, 2003 ).

2.2. Oxidation of bromide

2.2.1. Water treatmentIf bromide-containing water is treated with chlorine, bromideis oxidized to hypobromous acid and hypobromite (HOBr/

Br -

HOBr A

ABr

AO - + H +

H2O

Oxidant

Oxidant

i)

ii)

iii)

Fig. 1 e Role of bromide (Br L ) in oxidative water treatment:BrL is oxidized to bromine (HOBr for simplicity), whichreacts with organic or inorganic reaction partners (A),leading to (i) brominated compounds (ABr), (ii) AO L viahydrolysis of ABr or (iii) to other oxidized species (e.g.,DBPs, A0 ) after further oxidation of ABr. In reactions (ii) and(iii) bromide acts as a catalyst.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 216

http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030


3/28

OBr , Eqs. (1) and (2)) (Bousher et al., 1986; Farkas et al., 1949;Kumar and Margerum, 1987 ):

HOClþ Br / HOBrþ Cl k1 ¼ ð1:55 6:84Þ 103 M 1 s 1 (1)

ClO þ Br / BrO þ Cl k2 ¼ 9 10 4 M 1 s 1 (2)

In chloramination processes, bromide can also be oxidized(Eqs. (3)e (6)). The reaction is rst-order in monochloramine,bromide, andthe protonconcentration and itsmain product isassumed to be a mixed haloamine, NHBrCl ( Bousher et al.,1989; Gazda et al., 1993; Trofe et al., 1980 ).

NH2Cl þ Hþ $ NH3Clþ K3 ¼ 2:8 101 M 1 (3)

NH3Clþ þ Br / NH3Brþ þ Cl k4 ¼ ð5 6Þ 104 M 1 s 1 (4)

NH3Brþ þ NH2Cl/ NHBrClþ NH4þ fast (5)

This leads to the following overall reaction (Eq. (6), the rateconstant is based on the decrease of NH 2Cl, for thedecrease of Br /increase of NHBrCl the rate constant is half the value):

2NH2Cl þ Hþ þ Br / NHBrClþ Cl þ NH4þ

k6 ¼ ð2:8 0:3Þ 106 M 2 s 1 (6)

Bromine in bromochloramine is highly reactive ( Valentine,1986). Further species can be formed depending on pHand the Br/Cl ratio ( Bousher et al., 1989 ). The reaction of bro-mochloramine with monochloramine leads to the decompo-sition of both haloamines into N 2, Cl , Br and H þ . Bromidethus catalyzes the autodecomposition of monochloramine(Vikesland et al., 2001 ). However, in the presence of NOM, theaccelerated autodecomposition is reduced, as bromamine

reacts mainly with NOM ( Duirk and Valentine, 2007 ).The oxidation of bromide by chlorine dioxide (ClO 2) is very

slow (< 0.05 M 1 s 1 at pH 8) and can be neglected undertypical water treatment conditions ( Hoigné andBader,1994 ).Ithas been shown that peracetic acid is also able to oxidizebromide to HOBr ( Booth and Lester, 1995 ). However, this is aswell a very slow process and is probably not relevant forbrominated disinfection by-product formation ( Dell’Erbaet al., 2007).

The formation of HOBr from Br by ozone has been studiedin the context of bromate (BrO 3 ) formation ( Haag and Hoigné,1983; von Gunten, 2003; von Gunten and Hoigne ´, 1994).Second-orderrate constants forthe reaction betweenbromide

and ozone have been determined (Eq. (7)) (Haag and Hoigne´,1983; Haruta and Takeyama, 1981; Liu et al., 2001 ).

O3 þ Br / BrO þ O2 k7 ¼ ð1:60 2:58Þ 102 M 1 s 1 (7)

As shown in Fig. 2, HOBr plays an important role as anintermediate in the bromate formation mechanism during ozonation. Bromate has been classied as a potential carcino-gen and a drinking water standard/guideline of 10 mg L 1 hasbeen established worldwide ( European Union, 1998; UnitedStates Environmental Protection Agency, 2009; World HealthOrganization, 2011 ). Bromate is formed by a complex multi-step mechanism including ozone and hydroxyl radicals ( OH,Fig. 2) (von Gunten and Hoigne ´, 1994; von Sonntag and vonGunten, 2012 ).

HOBr can also be formed in advanced oxidation processes(AOPs) such as UV/H2O2 or O3 /H2O2, where hydroxyl radicalsare the only or the main oxidants, via the formation of thebromine atom (Br ) (Pinkernell and von Gunten, 2001; vonGunten and Oliveras, 1998 ). Hydrogen peroxide (H 2O2), whichis present in such processes reduces HOBr to bromide with ahigh second-order rate constant ( Table 3 , Section 3.2) andtherefore fully (UV/H 2O2) or partially (O 3 /H2O2) inhibits furtheroxidation of HOBr to bromate or its reaction with other com-pounds ( Symons and Zheng, 1997; von Gunten and Oliveras,1997).

In oxidative water treatment, bromide is thus oxidized bythe three commonly used oxidants chlorine, chloramine andozone. The oxidation of bromide by chlorine and ozone isrelatively fast, with hypochlorous acid reacting about tentimes faster with bromide than ozone, while the oxidation of bromide by chloramine is a signicantly slower process.

2.2.2. Biological systemsHypohalous acids are also formed in biological systems suchas the mammalian host defense system ( Thomas et al., 1995;Weiss et al., 1986 ). White blood cells are able to oxidize bro-mide (and chloride) with H 2O2. These reactions are catalyzedby myeloperoxidase (MPO) and eosinophil peroxidase (EPO),two heme enzymes ( Thomas et al., 1995 ). The hypohalousacids react with different biological molecules of pathogensbut can e if produced in excess or unintentionally e also harmthe tissue of the respective organism itself, leading to in-ammatory diseases ( Pattison and Davies, 2006 ).

2.2.3. Atmospheric watersBromine plays an important role in the troposphere at themarine boundary layer in a number of processes such asozone depletion, reactions with organic compounds, HO x andNOx , and oxidation of Hg(0) ( Finlayson-Pitts, 2003; Saiz-Lopezand von Glasow, 2012; Sommariva and von Glasow, 2012 ).The formation of bromine in the marine boundary is complex,involving reactions in the gas phase, in aerosols and on saltsurfaces such as the oxidation of bromide by ozone and HOCl(formed by the reaction of chlorine monoxide with hydro-peroxyl radical) ( Finlayson-Pitts, 2003; Vogt et al., 1996 ).

Fig. 2 e Bromate formation during ozonation of bromide-containing waters (from von Sonntag and von Gunten2012 , with permission).

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 17



4/28

Furthermore, similar to the discussion in Section 2.3, speciessuch as Br 2 and BrCl are also formed in aerosols and in ex-change with the gas phase.

2.3. Speciation of bromine in aqueous solution in theabsence and presence of chlorine

The reactions occurring in the Cl e HOCle Br system arecompiled in Table 1 . Figs.3 and 4 show the reactions/equilibriaof bromine species in the absence and presence of chlorine,respectively.

According to Fig. 3, HOBr undergoes acid e base speciation(Eq.(8), Table 1 ). Table 2 lists reported p Ka values,with 8.8 0.1being selected for low ionic strength.

HOBr disproportionates leading to the formation of bro-mide and bromate ( Fig. 3). The reaction occurs in two stepsaccording to Eqs. (25) and (26) (Br(-I): Br , Br(I): HOBr/BrO,Br(III): HBrO2 /BrO2 , Br(V): HBrO3 /BrO3 ) (Beckwith andMargerum, 1997 ).

BrðIÞ þ BrðIÞ/ BrðIIIÞ þ Brð IÞ k25 ¼ 2 10 3 M 1 s 1 (25)

BrðIÞ þ BrðIIIÞ/ BrðVÞ þ Brð IÞ faster than ð25Þ (26)

The reaction of HOBr with bromite (Eq. (26)) is fasterthan reaction (25), the latter thus dominating the

disproportionation ( Beckwith and Margerum, 1997 ). Theoverall disproportionation can be formulated as

3HOBr/ BrO3 þ 2Br þ 3Hþ (27)

Table 1 e Reactions of bromine and chlorine species: Equilibrium and rate constants. a

No. Reaction Equilibrium constant a Rate constant a

(1) HOCl þ Br $ HOBr þ Cl 1.5 105 k1 ¼ (1.55e 6.84) 103 M 1 s 1b

(8) HOBr$ OBr þ Hþ See Table 2(9) Br2 þ H2O$ HOBr þ Br þ Hþ (6.1 0.1) 10 9 M2c k9 ¼ 9.7 101 s 1d

k 9 ¼ (1.6 0.2) 1010 M 2 s 1d

(10) Br2 þ Br $ Br

3(1.61 0.03) 101 M 1e ,f k

10 ¼ 9.6 108 M 1 s 1g

k 10 ¼ 5.5 107 s 1g

(11) 2HOBr$ Br2O þ H2O 6.31 M 1h

(12) HOCl$ OCl þ Hþ pKa ¼ 7.47(13) Cl2 þ H2O$ HOCl þ Cl þ Hþ (1.04 0.07) 10 3 M2i k13 ¼ (2.23 0.06) 101 s 1i

k 13 ¼ (2.14 0.08) 104 M 2 s 1i

(14) BrCl þ H2O$ HOBr þ Cl þ Hþ (1.3 0.1) 10 4 M2 k14 ¼ (3.0 0.4) 106 s 1

(15) BrCl þ H2O$ HOCl þ Br þ Hþ 8.7 10 10 M2 k 15 ¼ 1.32 106 M 2 s 1

(16) 2BrCl$ Cl2 þ Br2 (7.6 1.7) 10 3

(17) BrCl þ Cl $ Cl2 þ Br 9.1 10 7

(18) Cl2 þ Cl $ Cl3 (1.8 0.2) 10 1 M 1e

(19) Cl2 þ Br $ BrCl2 4.2 106 M 1 k19 ¼ 7.7 109 M 1 s 1e

(20) BrCl þ Cl $ BrCl2 3.8 0.3 M 1 j

(21) Br2 þ Cl $ Br2Cl 1.3 0.3 M 1e

(22) BrCl þ Br $ Br2Cl (1.8 0.2) 104 M 1e k22 > 108 M 1 s 1e

(23) HOBr þ HOCl$

BrOClþ H2O 3.47 101

M1h

(24) 2HOCl$ Cl2O þ H2O 1.51 10 2 M 1h ,k

a Liu and Margerum (2001) and references within unless otherwise indicated, m¼ 1 M, 25 C.b Bousher et al. (1986), Farkas et al. (1949), Kumar and Margerum (1987) .c Beckwith et al. (1996) , m¼ 0.5 M, 25 C; Tó th and Fá bián (2004): (7.17 0.04) 10 9 M2, m¼ 1 M, 25 C; for an overview of K at various ionic

strengths and reported values, see Beckwith et al. (1996) .d Beckwith et al. (1996) , m¼ 0.5 M, 25 C; Eigen and Kustin (1962) : k9 ¼ 110 s 1, k 9 ¼ 1.6 1010 M 2 s 1, m¼ 0.1 M, 20 C.e Wang et al. (1994) , m¼ 1 M, 25 C.f Tóth and Fá bián (2000): (1.93 0.12) 101 M 1, m¼ 1 M, 25 C; Ershov (2004): 1.75 101 M 1.g Ershov (2004); other values reported: k10 ¼ (1.5 0.4) 109 M 1 s 1, k 10 ¼ (5 1.3) 107 s 1 (Ruasse et al., 1986 ).h Sivey et al. (2013).i Wang and Margerum (1994) , m ¼ 0.5 M, 25 C; an overview of reported values is also given there. j Wang et al. (1994) : 6.0 0.3 M 1, m¼ 1 M, 25 C.k Roth (1929): 8.71 10 3 M 1 at 19 C.

HOBr

Br 2

Br 3-

Br -, H +

Br -

HOBr/OBr - Br -

BrO 2 - BrO 3 -

Br 2O

(CuO)

HOBr/OBr - Br -

(CuO)

HOBr

OBr - + H +

Fig. 3 e Bromine species and their formation pathways inabsence of Cl L and HOCl: HOBr dissociates with a p K a of 8.8(see Table 2 ). At low pH, Br 2 is formed, which can react further to Br 3 L . HOBr may form Br 2 O in a self-reaction.Disproportionation of HOBr leads to Br L , BrO2 L and nallyBrO3 L . This is a slow process unless catalyzed by CuO orother corrosion products.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 218

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://dx.doi.org/10.1016/j.watres.2013.08.030http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


5/28

Disproportionation is fastest in the pH range 3 e 8, with amaximum at pH 7.3 ( Beckwith and Margerum, 1997; Chapin,1934). Above pH 8, disproportionation is signicantly slowerdue to the deprotonation of HOBr, while initial steps of thedisproportionation are thought to be reversible below pH 3,thus decreasing the rate of reaction ( Beckwith and Margerum,1997). Beckwith and Margerum (1997) gave an overview of reported rate constants at various pH values. It is suggestedthat the disproportionation is accelerated by bases, metalsand light ( Beckwith and Margerum, 1997 ). Under acidic con-ditions (pH < 4), molecular bromine is formed in addition tobromate in the disproportionation process (Eq. (28)) (Beckwithand Margerum, 1997 ).

5HOBr$ 2Br2 þ BrO3 þ 2H2O þ Hþ K28 ¼ 1 1010 M 1 (28)

Due to the low bromine concentration and the low rateconstant, disproportionation is a very slow process undertypical water treatment conditions and bromate formed bythis pathway is negligible. However, in case of concentratedhypochlorite solutions containing bromide, disproportion-ation becomes important and can lead to a dosing of bromateduring chlorination of drinking water ( Weinberg et al., 2003 ).

Furthermore, it has been observed that cupric oxide (CuO), acorrosion product of copper pipes, catalyzes the dispropor-tionation of hypobromous acid ( Liu et al., 2012). This may leadto elevated bromate concentrations during chlorination of bromide-containing waters if copper pipes are used in drink-ing water distribution systems ( Liu et al., 2012).

With the equilibria listed in Table 1 and illustrated in Fig. 4,the chlorine andbromine equilibrium-speciationas a functionof pH was calculated using PHREEQC ( Parkhurst and Appelo,1999) and based on a model by Korshin (2011). Threedifferent conditions have been considered: water in absenceof chlorine with 99% of bromide oxidized to bromine ( Fig. 5a),fresh water after chlorination with 2 mg L 1 total activechlorine ( Fig. 5b) and seawater after chlorination with2 mg L 1 total active chlorine ( Fig. 5c). The fresh water systemcontained 100 mg L 1 bromide and 10 mgL 1 chloride, whichrepresents average fresh water concentrations for both ions.In seawater, concentrations of 67 mgL 1 and 20 gL 1 wereassumed for bromide and chloride, respectively. Seawater isoxidized/disinfected, e.g., for treatment of ballast water, aquacultures or seawater aquaria ( Gonçalves and Gagnon, 2011;Grguric and Coston, 1998; Jorquera et al., 2002; Werschkunet al., 2012). In absence of chlorine ( Fig. 5a), HOBr and OBrdominate the system. Br 2 and (to a much lesser extend) Br 3gain importance under acidic conditions with their concen-trations depending on the residual bromide concentration.The concentration of Br 2O is proportional to [HOBr] 2. In chlo-rinated fresh water systems ( Fig. 5b), HOBr and OBr remainthe dominant bromine species. In the pH range relevant forwater treatment (pH 6 e 8), the concentrations of other speciessuch as BrCl, BrOCl, Br2O and Br2 are ve to eight orders of magnitude lower than the HOBr concentration. Due to theirlow concentrations, their reactivity with compounds thatreact quickly with HOBr can normally be neglected underwater treatment conditions, even though these species canalso be very reactive ( Odeh et al., 2004; Sivey et al., 2013 ).BrCl2 , Br3 and Br 2Cl are not relevant under the conditionsconsidered here. In the case of chlorinated seawater ( Fig. 5c),

Br - HOBr

Cl -HOCl

HOBrHOCl

Br 2OBrOCl

Br 2 Br 3-Br

-Cl2Cl3-

Cl -

Cl -, H +

Br -, H +

BrCl 2- Br 2Cl-

Cl -Br -

Br -Cl -

OBr - + H +

OCl - + H +

HOCl

Cl2O

BrCl

Fig. 4 e Overview of the formation pathways and the speciation of all chlorine, bromine and brominechlorine species in theClL e HOCle BrL system. Equilibria and rate constants are compiled in Table 1 . HOBr and HOCl disproportionations have been neglected since they are only relevant at high concentrations or for very long reaction times.

Table 2 e Dissociation constants for HOBr.

ReportedpKa

Ionicstrength

[M]

Temperature[ C]

Reference

8.76 0.02e 0.15 20 (Haag and Hoigne´, 1983)8.8 ± 0.1 * 0.5 25 (Troy and Margerum, 1991 )8.59 1 25 (Gerritsen et al., 1993 )7.87 1 25 (Lahoutifard et al., 2002 )8.7 n.a. 20 (Shilov, 1938 )8.7 n.a. n.a. (Farkas and Lewin, 1950 )

n.a. ¼ not available.* Selected value.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 19



6/28

the concentrations of most bromine species arehigher than infresh water due the higher natural abundance of chloride andbromide. HOBr and OBr are still the most important brominespecies in the pH range 6 e 8. Yet, other bromine species(i.e. Br2, Br2Cl , BrCl, Br3 , BrCl2 and Br 2O) gain importanceand their concentrations are only about two to four orders of magnitude lower than the HOBr concentration at pH 7. AtpH < 5, Br2 and Br2Cl become the dominant species.

3. Reactions of HOBr/OBr L with inorganicand organic compounds

3.1. Kinetic considerations

In most cases, the oxidation of 1 mol of a compound C by h

moles of oxidant Ox (e.g., bromine) resulting in 1 mol of a

-15

-13

-11-9

-7

-5

-3

-15

-13

-11

-9

-7

-5

-3

-15

-13

-11

-9

-7

-5

-3

4 5 6 7 8 9 10 11 12

pH

l o g

( c o n c e n

t r a

t i o n ,

M )

a)

c)

b)

Fig. 5 e Concentrations of halogen species as a function of pH in (a) fresh water containing 100 m g LL 1 bromide( [ 1.25 3 10 L 6 M), of which 99% is oxidized to bromine, without active chlorine, (b) fresh water containing 100 m g LL 1

bromide ( [ 1.25 3 10 L 6 M) and 10 mg L L 1 chloride ( [ 2.82 3 10L 4 M), which is chlorinated with 2 mg L L 1 ( [ 2.82 3 10L 5 M)active chlorine, (c) seawater containing 67 mg L L 1 bromide ( [ 8.38 3 10L 4 M) and 20 g L L 1 chloride ( [ 0.56 M), which ischlorinated with 2 mg L L 1 ( [ 2.82 3 10L 5 M) active chlorine. Calculations were done using PHREEQC ( Parkhurst and Appelo,1999 ), based on the equilibrium constants in Table 1 . The shaded area represents the typical pH range for water treatment.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 220



7/28

product P can be described by second-order kinetics, rst-order with respect to each reaction partner ( von Sonntag and von Gunten, 2012 ):

1h

d½Oxdt ¼

d½Cdt ¼

d½Pdt ¼ kox;C ½Ox ½C (29)

where kox, C is the second-order rate constant [M 1 s 1].

In case compound C and/or the oxidant undergo acid e baseequilibria, the speciation has to be taken into account. If hypobromous acid reacts with a compound that does notspeciate, the oxidation of the compound by the protonatedform (HOBr) and the anion (OBr ) has to be considered in therate law, which leads to

d½Cdt ¼ kHOBr;C ½HOBr ½C þkOBr ;C OBr ½C (30)

with kHOBr, C and kOBr ;C being the species-specic rate con-stants. Based on the pH and the dissociation constant Ka , thedegree of dissociation a HOBr (i.e. the protonated fraction) canbe dened (Eq. (31)).

a HOBr ¼ ½Hþ

½Ka þ ½Hþ (31)

Eq. (30) then becomes:

d½Cdt ¼ kHOBr;C

a HOBr þ kOBr ;C ð1 a HOBrÞ ½HOBr tot ½C

(32)

The expression in brackets represents the apparent second-order rate constant kapp , which is pH-dependent:

kapp ¼ kHOBr;C a HOBr þ kOBr ;C ð1 a HOBrÞ (33)

If compound C speciates as well (as CH/ C ), the rate lawand the apparent rate constant have to be extended

accordingly (Eq. (34), (Criquet et al., 2012; von Gunten andOliveras, 1997 )).

kapp ¼ kHOBr;C a HOBr ð1 a CHÞ þ kOBr ;CH ð1 a HOBrÞ a CHþ kHOBr;CH a HOBr a CH þ kOBr ;C ð1 a HOBrÞ

ð1 a CHÞ(34)

In the following sections, rate constants for the reaction of bromine with various classes of compounds are discussed.Besides the kinetics for the reactions with HOBr andOBr , rateconstants at lower pH, indicated by HOBr þ Hþ , have oftenbeen proposed in literature. It has to be noted that these re-actions may also reect the reactivity of Br 2, which is formedunder these conditions (Eq. (9), Fig. 5a). Furthermore, incertain cases rate constants for Br 2 and Br3 are also given.

3.2. Reactions of bromine with inorganic compounds

Table 3 provides a compilation of rate constants for the re-action of HOBr and OBr with various inorganic compounds.

3.2.1. AmmoniaThe reaction of HOBr with ammonia to NH 2Br is very fast (Eq.(35)).

HOBrþ NH3$ NH2Br þ H2O k35 ¼ ð4 7:5Þ 107 M 1 s 1 (35)

Judging from two conictingequilibrium constants that havebeen reported ( Soulard et al., 1981; Tremblay-Goutaudier et al.,1994), the rate constant for the back reaction should be in theorder of 1 e 103 s 1. In a model, a value of 1.5 10 3 s 1 wasestimated ( Haagand Lietzke,1980 ),while a tted hydrolysis rate

constant of 7.5 106 M 1 s 1 which is rst-order in OH wasused in another model by Pinkernell and von Gunten (2001) .

Table 3 e Rate constants for the reaction of inorganic compounds with bromine.

Compound p Ka Species-specic rate constants Apparent rateconstant at pH 7

Temperature Reference

kðHOBrþ Hþ ÞM 2 s 1

k(HOBr)[M 1 s 1]

kðOBr ÞM 1 s 1

kapp[M 1 s 1] [ C]

Br See Table 1BrO 8.8 Disproportionation:

See Section 2.3ClO2 2.0 9.7 101 9.6 101 25 (Furman and Margerum, 1998 )CN 9.2 (4.2 0.9) 109 (5.7 0.4) 107 2.6 107 25 (Gerritsen et al., 1993 )HO2 11.6 (7.6 1.3) 108 1.9 104 25 (von Gunten and Oliveras, 1997 )I 10 (5.0 0.3) 109 (6.8 0.4) 105 4.9 109 25 (Troy and Margerum, 1991 )OI 1.9 106 1.8 103 7.3 102 24 1 (Criquet et al., 2012 )HOI 10.4 Not signicant 24 1 (Criquet et al., 2012 )IO2 High High 24 1 (Criquet et al., 2012 )NH3 9.3 (7.5 0.4) 107 (7.6 0.4) 104 4.1 105 20 (Wajon and Morris, 1982 )

(4 1) 107 2.2 105 25 (Inman and Johnson, 1984 )NH2Cl 1.4 (2.9 0.1) 105 (2.2 0.1) 104 2.8 105 25 (Gazda and Margerum, 1994 )NO2 3.4 (1.4 1) 104 1.4 104 (Lahoutifard et al., 2002 )O2 4.8 3.5 109 3.4 109 (Schwarz and Bielski, 1986 )SCN 1.1 2.3 109 3.8 104 2.3 109 18 (Nagy et al., 2006)SO32 7.2 (5 1) 109 (1.0 0.1) 108 1.9 109 25 (Troy and Margerum, 1991 )

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 21



8/28

The second-order rate constant for the reaction of HOBrwith NH 3 is approximately three orders of magnitude higherthan the corresponding rate constant forOBr (Table 3 ), whichis due to the higher electrophilicity of HOBr compared to OBr .The activation energies for these reactions are 15.7 (HOBr)and48.4 kJ mol 1 (OBr ) (Wajon and Morris, 1982 ). The mechanismof the reaction betweenHOBr and ammonia is an electrophilic

substitution of H(I) in ammonia by Br(I), forming monobrom-amine ( Wajon and Morris, 1982 ). For comparison, the species-specic rate constant for the reaction of HOCl with ammoniais about one order of magnitude smaller in the order of 106 M 1 s 1 (Deborde and von Gunten, 2008 ).

Fig. 6 shows the apparent second-order rate constants forthe reaction of chlorine and bromine with ammonia as afunction of pH. Since the main reacting species are NH 3 andHOBr or HOCl, respectively, the pH of the maximum apparentsecond-order rate constants equals the average of the p Kavalues of NH 4 þ and the hypohalous acid (pH 9.1 for HOBr/NH 3and 8.4 for HOCl/NH3). At pH 7, the difference between theapparent second-order rate constants of the two reactions(HOBrþ NH3 vs. HOClþ NH3) is a factor of around 30, at pH 8 afactor of around 80. The fraction of ammonia ð f HOBr;NH3 Þreacting with HOBr at a certain pH can be calculated by Eq.(36):

f HOBr;NH3 ¼ kapp ;HOBr ½HOBr tot

kapp ;HOBr ½HOBr tot þ kapp ;HOCl ½HOCl tot(36)

with kapp,HOBr and kapp,HOCl being the apparent second-orderrate constants for the reaction of HOBr and HOCl withammonia. Under the conditions in Fig. 5b (water containing 1.25 10 6 M bromide, chlorinated with 2.8 10 5 M activechlorine) and assuming that bromide was oxidized to HOBr

prior to ammonia addition, at pH 7 around half of theammonia present reacts with HOBr, while at pH 8, thefractionincreases to around 80%. However, in the more realistic caseof a water containing both bromide and ammonia at the sametime when chlorinated, the majority of chlorine will reactimmediately with ammonia (fraction of ammonia reacting with chlorine > 99%, assuming the same concentrations as

beforeandwith [NH 3]tot ¼ [HOCl]tot at neutral pH), suppressing the oxidation of bromide in the rst place. Only if theammonia concentration is lower than the active chlorineconcentration, a relevant bromine formation will be observed.The reactions of the formed chloramine with bromide arediscussed above (Section 2.2.1).

Due to its high reactivitywith HOBr, ammonia addition has

been suggested as a measure for bromate mitigation during ozonation of bromide-containing waters, because ammoniamasks HOBr, a decisive intermediate species in the bromateformation pathway ( Haag et al., 1984; Pinkernell and vonGunten, 2001 ). Based on the masking of HOBr as NH 2Br, thechlorine e ammonia process has been developed ( Neemannet al., 2004). In this process, a pre-chlorination to oxidizebromide to HOBr is carried out, followed by the addition of ammonia to mask HOBr as NH 2Br before ozone is added.Compared to conventional ozonation, bromate formation canbe reduced by up to a factor of ten by this process ( Bufe et al.,2004). This process can be applied if high ozone exposures arerequired (e.g., for Cryptosporidium parvum oocysts inactivation)in waters with elevated bromide levels ( Bufe et al., 2004).

Bromine can react further with NH 2Br to form dibrom-amine (NHBr 2) and eventually tribromamine (NBr 3) (Galal-Gorchev and Morris, 1965 ). The relative concentrations of the three bromamine species after reaction depend on the pHand the N/Br ratio ( Galal-Gorchev and Morris, 1965; Johnsonand Overby, 1971 ). Dibromamine is also formed through thedisproportionation of monobromamine to dibromamine andammonia (Eq. (37)) and analogously tribromamine can beformed ( Inman and Johnson, 1984; Lei et al., 2004; Soulardet al., 1981).

2NH2Brþ Hþ $ NHBr2 þ NH4þ K37 ¼8 107 5 109 M 1 (37)

It was reported that the disproportionation of monobrom-amine/formation of dibromamine is catalyzed by phosphatebuffer and the following rate law for the pH range 7 e 8.5 wasproposed ( Inman and Johnson, 1984 ) (Eq. (38)):

d½NHBr2 =dt ¼ k38a Hþ ½PO4 ½Br 2tot þ k38b Hþ ½Br2tot (38)

[Br]tot : sum of HOBr and OBr , [PO4]: total phosphate concen-tration; with k38 a ¼ 9.9 1011 M 3 s 1 and k38 b ¼ 2.4 108 M 2

s 1 (Inman and Johnson, 1984 ). The catalytic effect of acids onthe disproportionation of monobromamine was conrmed,however with different rate constants in another study ( Leiet al., 2004). The measured rate constants for the forward re-action ranged from 0.5 M 1 s 1 (with H 2O as a catalyst) to5 108 M 2 s 1 (Hþ catalysis) and for the backward reactionrate constants varied between 1 (with H 2O as a catalyst) and1 109 M 2 s 1 (Hþ catalysis).

While the disproportionation of monobromamine/forma-tion of dibromamine is a reversible process (Eq. (37)), furtherdecomposition reactions of monobromamine and dibrom-amine (Eqs. (39) and (40)) were found to be irreversible andbase-catalyzed ( Lei et al., 2004).

NH2Br þ NHBr2/ products (39)

2NHBr2/ products (40)

0.0

0.5

1.0

1

2

3

4

5

6

7

8

4 6 8 10 12

R e

l a t i v e s p e c

i e s c o n c e n

t r a

t i o n

l o g

( k a p p ,

M - 1 s -

1 )

pH

k a p p b r o

m i n e

k a p p c h

l o r i n e

HOBr HOCl

NH3

Fig. 6 e Reaction of HOX with ammonia: pH dependence of reactive species concentration and apparent second-orderrate constants for the reaction of HOBr and HOCl withammonia.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 222



9/28

The corresponding rate constants were determined in H 2Oand for catalysis by OH , PO4 3 , CO3 2 , NH3 and HPO4 2 . Therate constants without catalysis are 8.9 M 1 s 1 (Eq. (39)) and6.2 M 1 s 1 (Eq. (40)); the values for OH catalysis are4.1 107 M 2 s 1 (Eq. (39)) and 8.3 104 M 2 s 1 (Eq. (40)), therate constants for catalysis with other bases lie between thetwo values for H 2O and OH (Lei et al., 2004).

Little is known about the rate and mechanism of tribrom-amine formation ( Hofmann and Andrews, 2001 ); its decom-position is discussed by Inman et al. (1976) and Johnson andOverby (1971).

3.2.2. Reactivity of bromaminesNot much information is available on the reactivity of brom-amines with inorganic andorganic compounds. The reactivityof bromine (Br(I)) with NOM in a mixed bromide e chloraminesystem was investigated by Duirk and Valentine (2007) .Another study discusses the kinetics of the reaction betweenmono- and dibromamine and CN (Eqs. (41) and (42)). Theproduct of these reactions is in both cases cyanogen bromide(BrCN) (Lei et al., 2006).

NH2Br þ CN þ H2O/ NH3 þ BrCNþ OHk41 ¼ 2:63 104 M 1 s 1

(41)

NHBr2 þ CN þ H2O/ NH2Br þ BrCNþ OHk42 ¼ 1:31 108 M 1 s 1

(42)

The rate constants of Eqs. (41) and (42) are ve to six orders of magnitude higher than the analogous rate constants of chlo-ramines ( Lei et al., 2006). Compared to HOBr, the second-orderrate constant for the reaction of NHBr 2 with CN is only a

factor of about 30 lower ( Table 3 ). The reaction of mono-bromamine with cyanide is general-acid-catalyzed, while nocatalytic effect was found for the reaction with dibromamine(Lei et al., 2006). Third-order catalysis rate constants for H þ ,H2PO4 , HPO4 2 , H3BO3 and NH4 þ have also been determined(Lei et al., 2006). For the hydrolysis of BrCN, see Section 3.2.3.

The reactions of bromamines with acetic acid forming haloacetic acids were studied by Pope and Speitel (2008) . Itwas found that bromamines are signicantly more reactivewith acetic acid than their chlorine analogs ( Pope and Speitel,2008).

The reaction of monobromamines with ozone yields bro-mide and nitrate (Eq. (43)) (Haag et al., 1984 ).

NH2Br þ 3O3/ Br þ NO3 þ 3O2 þ 2Hþ

k43 ¼ ð4:0 1:0Þ 101 M 1 s 1 (43)

For the analogous reaction of dibromamine with ozone, thesecond-order rate constant is a factor of four smaller ( Haag et al., 1984). No rate constant has been determined for tri-bromamine. Because bromamines react faster with ozonethan ammonia/ammonium at circumneutralpH, an enhancedammonia oxidation during ozonation of bromide-containing waters has been proposed ( Haag et al., 1984 ).

The reaction of monobromamine with biological sub-strates has been studied in the context of inammatory dis-eases. The apparent second-order rate constant of NH 2Br wasfound to be 1.8 103 M 1 s 1 (pH 7.3) for methionine,

> 105 M 1 s 1 (pH 7.2) for glutathione, 3.6 103 M 1 s 1 forascorbate (pH 6.9) and 1.02 102 M 1 s 1 for Fe(III)cytochromec (at pH 7.2) (Pru ¨ tz et al., 2001 ). Under acidic conditions, thereaction of monobromamine with iodide forms IBr andammonia ( Pru ¨ tz et al., 2001 ). For the reactivity of organicbromamines, see Section 3.3.2.

3.2.3. Halides and other inorganic anionsBased on the equilibrium constant and the rate constant forthe reaction of HOCl with Br , it can be assumed that the backreaction, i.e. the reaction of HOBr with Cl , is quite slow with asecond-order rate constant in the order of 10 2 M 1 s 1 (seeTable 1 ). As far as the reaction of HOBr with the other halidesis concerned, HOBr reacts with Br to form Br 2 in acidic so-lution (Eq. (9), Table 1 ) and oxidizes I in a very fast process(Eq. (44), Table 3 ) leading to the formation of the intermediateIBr. IBr hydrolyzes quickly to hypoiodous acid (HOI/OI ) withhydrolysis constants of 8 105 s 1 with H 2O and6 109 M 1 s 1 with OH (Troy et al., 1991 ).

HOBrþ I / HOIþ Br k44 ¼ ð5:0 0:3Þ 109

M1

s1

(44)HOI/OI can be further oxidized to iodite (IO 2 , Eq. (45)), whichis then again quickly oxidized to iodate (IO 3 , Eq. (46)) by HOBr(Criquet et al., 2012 ).

HOBrþ OI / IO2 þ Br þ Hþ k45 ¼ 1:9 106 M 1 s 1 (45)

HOBrþ IO2 / IO3 þ Br þ Hþ faster than ð45Þ (46)

The oxidation of I to iodate by HOBr is preferable in drinking water treatment because iodate is non-toxic ( Bu ¨ rgi et al., 2001)and concurrently the formation of iodo-organic DBPs isminimized ( Criquet et al., 2012 ). In the sequence of iodide

oxidation to iodate (Eqs. (44)e

(46)), the oxidation of hypoiodite(Eq. (45)) is the rate-determining step ( Criquet et al., 2012 ). Therate constant for the reaction of HOBr with HOI is considerablylower than the one with OI , because of the higher nucleo-philicity of OI compared to HOI. Analogous reactions alsooccur with OBr , however, the corresponding second-orderrate constants are about three to four orders of magnitudelower (Table 3 ).

Similar to the reaction of HOBr with iodide, reactions of HOBr with sulte (SO3 2 ), cyanide (CN ) and thiocyanate(SCN ) are also all very rapid and close to the diffusion limit(Gerritsen et al., 1993; Nagy et al., 2006; Troy and Margerum,1991). Typically, the reaction of HOBr is one to nearly ve or-

ders of magnitude faster than OBr (Table 3 ). The proposedpathway for the reactions with I , SO3 2 and CN is a Brþ

transfer resulting in IBr, SO 3Br and CNBr as intermediates(Fig. 1, (i)), which subsequently hydrolyze to HOI, SO 4 2 andOCN , respectively (Eq. (47), Fig. 1, (ii)) (Gerritsen et al., 1993;Troy and Margerum, 1991 ).

HOBrþ A $ ABrþ OH $ AOHe

AO þ Hþþ Br (47)

SO3Br hydrolyzes to SO 4 2 with a rst-order rate constantof 230 s 1 (0 C) (Troy and Margerum, 1991 ) and for CNBr hy-drolysis rate constants of 0.53 M 1 s 1 with OH and7.5 10 3 M 1 s 1 with CO3 2 were reported, resulting in the

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 23



10/28

formation of OCN and Br (Gerritsen et al., 1993 ). For the fastreaction of SCN with HOBr two parallel pathways of general-acid-catalyzed Br þ transfer and direct reaction of HOBr wereproposed. The product of these reactions is OSCN (Nagyet al., 2006).

Species-specic second-order rate constants for the re-actions ofbromine with the anions discussed above (I , SO3 2 ,

CN , SCN ) are 2e 100 times higher than those of chlorineðkðHOCl;I Þ ¼ 1:4 108; kðHOCl;SO 23 Þ ¼ 7:6 10

8; kðHOCl;CN Þ ¼1:2 109 and kðHOCl;SCN Þ ¼ 2:3 107 M 1 s 1Þ (Ashby et al.,2004; Deborde and von Gunten, 2008 ). The differences aremore pronounced (up to a factor of around 350) whencomparing the respective apparent second-order rate con-stants at pH 7 and 8 ( Table 4 ).

For nitrite, a rate constant of 1.4 104 M 1 s 1 in 1 MNa2SO4 was reported by Lahoutifard et al. (2002) . This rateconstant was about a factor of 3 smaller when the ionicstrength was adjusted with NaCl instead of Na 2SO4(Lahoutifard et al. 2002 ). However, the bromine speciationchanges signicantly under these conditions (see Fig. 4,Table 1 ). The reaction pathway proposed by Lahoutifard et al.(2002) consists of a fast formation of BrNO 2 (nucleophilicattack of Br þ ), which then e in the rate-determining step ereacts with NO 2 to N2O4. N2O4 hydrolyzes to NO 2 and NO 3(Lahoutifard et al., 2002 ). This partly agrees with a mechanismsuggested for the oxidation of NO 2 by HOCl ( Johnson andMargerum, 1991 ). After the formation of ClNO 2 two path-ways were observed: either N 2O4 formation and subsequenthydrolysis or decay of ClNO 2 yielding Cl and NO 2 þ , whereNO2 þ hydrolyzes to NO 3 ( Johnson and Margerum, 1991 ).

The rate constants for the reaction of HOBr with HO 2 (vonGunten and Oliveras, 1997 ) and O2 (Schwarz and Bielski,1986) are very high ( Table 3 ). The reaction with HO 2 is of particular interest in H 2O2-based AOPs (i.e. O3 /H2O2, UV/H2O2)as it minimizes the formation of bromate and brominatedDBPs as discussed in Section 2.2.1.

3.2.4. Iron, manganese and arsenicThe oxidation of Fe(II) and Mn(II) by bromine has not receivedmuch attention in literature except in the context of chemicaloscillation reactions (e.g., Chou et al., 1993; Hegedu ¨ s et al.,2006; Melichova et al., 1995, 2001; Sasaki, 1990 ). Also for the

oxidation of arsenite (As(III)) by bromine, no rate constant isavailable in literature.

In general, it can be assumed that an increase in pH willresult in a faster reaction of Fe(II) due to the increasing con-centration of Fe(II) hydroxy complexes, which are more sus-ceptible to oxidation ( Deborde and von Gunten, 2008; King,1998).

In the case of manganese it is known that, besides the slowdirect oxidation of Mn(II) by bromine, mostly heterogeneousMn(III,IV) oxide-catalyzed reactions occur: Mn(II) adsorbs toMn(III,IV) oxide, which results in a faster oxidation of Mn(II)(Allard et al., 2013; Melichova et al., 2001 ). This autocatalyticmechanism has also been observed for the analogous Mn(II)oxidation by chlorine ( Hao et al., 1991 in Deborde and vonGunten, 2008 ). It has been suggested that the oxidation of Mn(II) by bromine occurs to a signicant extent through BrCland Br 2O (Allard et al., 2013). This is thus a case wherebrominated electrophiles other than HOBr need to be takeninto account (see Section 2.3).

Because no rate constants for the reaction of Fe(II) andAs(III) with bromine are available, the corresponding rateconstants for chlorine may give an indication (Fe(II):1.7 104 M 1 s 1 at acidic pH ( Folkes et al., 1995); As(III):4.3 103 M 1 s 1 for As(OH)3, 5.8 107 M 1 s 1 for As(OH)2Oand 1.4 109 M 1 s 1 for As(OH)O2 2 (Dodd et al., 2006 )). Basedon these rate constants, it can be expected that the corre-sponding rate constants for the reaction of bromine with Fe(II)and As(III) are very high.

3.3. Reactions of bromine with organic compounds

3.3.1. Aromatic compoundsTable 5 summarizes published rate constants for the reactionof bromine with various aromatic compounds. Most rateconstants are available for the reaction of bromine with phe-nols, since an important class of taste and odor compounds isattributed to the products of these reactions (bromophenols,medicinal taste and odor) ( Piriou et al., 2007 ). The odorthreshold concentrations are very low with 30 ng L 1 and0.5 ng L 1 for 2-bromophenol and 2,6-bromophenol, respec-tively (Acero et al., 2005 ).

The rate constants for the reaction of bromine with phenoland phenolic compounds are very high ( Table 5 ). Yet, thespecies-specic rate constants for the reaction with phenoldetermined by Gallard et al. (2003) and the phenolic com-pounds determined by Guo and Lin (2009) seem to be slightlyoverestimated due to the use of quench-ow systems, whichdo not allow proper mixing for such fast reactions. In thecompilation in Table 5 , a value based on new measurementswith an advanced quench-ow system has been added(Criquet et al., in preparation ).

Due to the acid/base speciation of both HOBr and phenol,the apparent reaction rate constant varies with pH, similar tothat observed for the reaction of HOBr with NH 3 shown above(Eq. (35), Table 3 , Fig. 6). A comparison of the species-specicrate constants clearly shows that the protonated form,i.e. HOBr has the highest reactivity with the phenolate ion(Table 5 ). This can be explained by the higher electrophilicityof HOBr compared to OBr and the increased electron density

Table 4 e Comparison of apparent second-order rateconstants for the reaction of chlorine and bromine withselected anions at pH 7 and pH 8 (for references of bromine rate constants see Table 3 ; chlorine rateconstants from Ashby et al., 2004 and Deborde and vonGunten, 2008 ).Compound p Ka kapp [M 1 s 1]

Apparent rateconstant at pH 7

Apparent rateconstant at pH 8

Chlorine Bromine Chlorine BromineCN 9.2 5.8 106 2.6 107 1.7 107 2.2 108

I 10 1.1 108 4.9 109 3.4 107 4.3 109

SCN 1.1 1.8 107 2.3 109 5.5 106 2.0 109

SO3 2 7.2 2.2 108 1.9 109 1.6 108 3.7 109

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 224



11/28

Table 5 e Rate constants for the reaction of aromatic compounds with bromine.

Compound p Ka Species-specic rate constants Apparent rate constantat pH 7 or at given pH

Tempe

kðBr2=PhOHÞM 1 s 1

kðBr2=PhO ÞM 1 s 1

k(HOBr/PhOH)[M 1 s 1]

kðHOBr=PhO ÞM 1 s 1

kðOBr =PhO ÞM 1 s 1

kapp[M 1 s 1] [

4-Acetylphenol 8.6 (1.0 0.4) 104 (4.1 0.6) 106 < 2.0 103 1.1 105 23N-Acetyl-tyrosine (2.6 0.2) 105 (pH 7.2e 7.5) 22

(5.8 0.2) 105 (pH 7.2e 7.5) 37

p-Aminophenol 10.3 8.9 104 5.4 108 4.9 103 3.5 105 25 Anisole 5.2 101 20 Benzene < 1 10 2 (pH 4) 20

1.1 10 5 25 2-Bromo-4-methylphenol (6.4 0.3) 105 (pH 7.2e 7.5) 222-Bromophenol 8.5 6.4 106 2.2 105 25

1.0 104 6.2 109a 25 4-Bromophenol 9.2 4.8 106 3.2 104 25

3.9 103 5.5 109 25 2-Chlorophenol 8.6 7.2 106 1.9 105 253-Chlorophenol 7.9 b 6.5 104 7.9 106 7.5 102 1.0 106 25 4-Chlorophenol 9.4 (6.0 3.0) 103 (7.0 0.8) 106 (5.5 4.0) 104 3.2 104 23m-Cresol (3-methylphenol) 10.1 (3.5 0.3) 108 2.7 105 20

o-Cresol (2-methylphenol) 10.3 (9.6 0.4) 107 4.5 104 20 p-Cresol (4-methylphenol) 10.3 < 1.0 104 (2.1 0.5) 108 < 3.0 105 1.2 105 23

(2.6 0.1) 105 (pH 7.2e 7.5) 22Cyclo(Serine-Tyrosine) w 1.2 105 (pH 7.8) 202,4-Dibromophenol 7.8 1.2 104 8.9 105 4.2 104 1.3 105 25

w 3.0 102 3.7 109 25 2,6-Dibromophenol 6.7 1.7 104 4.8 105 8.2 104 3.3 105 25

w 5.0 102 2.7 109 25 2,4-Dichlorophenol 7.9 (3.0 1.0) 104 (8.8 0.9) 105 (3.0 2.0) 104 1.3 105 232,6-Dichlorophenol 7.0 3.5 104 4.5 105 1.6 104 2.5 105 25Hydroquinone (2.4 0.2) 105 (pH 7.4) 224-Hydroxybenzoic acid 9.5 (1.4 0.1) 107 5.2 104 20 3-(4-Hydroxyphenyl)

propionic acid (HPPA)(1.6 0.3) 105 (pH 7.2e 7.5) 22

3-Methoxyphenol 9.7 6.8 105 6.5 108 7.6 104 2.1 106 25 4-Methoxyphenol 10.2 (4.9 0.6) 107 3.1 104 20 4-Nitro-3-methylphenol 7.5 105 3.2 108 5.6 102 25 4-Nitrophenol 7.2 9.2 103 8.8 106 6.4 103 3.6 106 25

< 6.0 101 1.2 109a 25

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


12/28

Table 5 e ( continued )

Compound p Ka Species-specic rate constants Apparent rate constantat pH 7 or at given pH

Tempe

kðBr2=PhOHÞM 1 s 1

kðBr2=PhO ÞM 1 s 1

k(HOBr/PhOH)[M 1 s 1]

kðHOBr=PhO ÞM 1 s 1

kðOBr =PhO ÞM 1 s 1

kapp[M 1 s 1] [

Phenol 10.0 < 5.0 102 (1.8 0.2) 108 < 1.0 105 1.8 105 23(4.1 0.1) 107 4.1 104 20

8.0 103 2.3 108 5.3 102 2.3 105 25 5.0 102 (pH 4) 20

< 1 6.6 107 3.5 104 6.5 104 25 4.3 105 w 1.2 109c 25

3-Phenylpropionic acid Very slow (pH 7.2 e 7.5) 22Pyrene 1.1 101 20 2,4,6-Tribromophenol 6.8 2.0 102 3.3 103 < 1 2.1 103 25

5.9 w 1.7 103a 25 2,4,6-Trichlorphenol 6.2 (1.4 0.1) 103 1 1.2 103 23Trolox (6.4 0.6) 104 (pH 7.4) 22Ubiquinol-0 (2.5 0.3) 106 (pH 7.4) 22Vanillin 7.4 (1.3 0.1) 106 6.0 105 20

p-Xylene 2 10 1 (pH 4) 201.1 10 4 20

a These values may be overestimated due to polybromination ( Tee et al., 1989 ).b pKa used by Guo and Lin (2009) for the tting of the rate constants. The CRC Handbook of Chemistry and Physics gives a value of 9.1 ( Haynes, 2013 ).c A second-order rate constant of 8.5 108 M 1 s 1 has been determined for the reaction of Br 3 with PhO (Tee et al.,1989 ). However, this is not relevant undetext).

http://-/?-http://-/?-http://-/?-http://-/?-


13/28

at the ortho and para position of the phenolatecompared to thephenol.

The mechanism of the reaction of bromine with phenols ismainly an electrophilic substitution caused by the positivepartial charge ( dþ ) of the Br in HOBr (Acero et al., 2005). Thebromination occurs in ortho or para position. Statistically, theortho position is favored which leads to approximately 2/3ortho substitution and 1/3 para substitution ( Acero et al., 2005 ).When all the ortho and para positionshave been substituted bya halogen, further halogenation leads to ring opening andformation of THMs ( Acero et al., 2005; Arnold et al., 2008;Gallard and von Gunten, 2002 ).

While electrophilic substitution is the main reactionmechanism of HOBr with phenolic compounds, for catecholand hydroquinone electron transfer has been observed, whichleads to the formation of 1,2-benzoquinone and 1,4-benzoquinone, respectively ( Criquet et al., in preparation ).

In addition to the rate constants of HOBr and OBr , second-order rate constants for the reaction of Br 2 with phenols andphenolates (and some other compounds, not listed) have beenmeasured ( Tee and Bennett, 1988; Tee and Iyengar, 1985; Teeet al., 1985, 1989) (Table 5 ). The second-order rate constantsfor the reaction of Br 2 with phenol and phenolates areextremely high, close to diffusion limitation. Furthermore, asecond-order rate constant has been determined for the re-action of Br 3 with phenolate. This is, however, irrelevantunder any natural water conditions, as very high bromideconcentrations would be necessary for Br 3 formation at a pHfor which there is sufcient overlap with phenolate (4 mM of bromide at pH 6 for Br 3 to contribute to 1% of the total re-action). In contrast, the reaction of Br 2 with phenol can play arole at low pH: In the case of a water initially containing 1.25 mM bromide (100 mg L 1) of which 90% has been oxidized(i.e. 1.25 10 7 M bromide remaining) at pH 6, around 1 & of the total phenol reacts with Br 2, while at pH 5 it is alreadyaround 10%.

The analogous reactions of HOCl with phenol and halo-phenols follow the same reaction pathway as observed forHOBr (Deborde and von Gunten, 2008 ). Fig. 7 shows the cor-relation between second-order rate constants for the reactionof HOCl ðkðHOCl;PhO ÞÞ and HOBr ðkðHOBr;PhO ÞÞ withvarious phenolates ( R2 ¼ 0.998). The slope of the linear corre-lation is 3000, which represents the ratio betweenðkðHOBr;PhO ÞÞ and ðkðHOCl;PhO ÞÞ. This means that the species-specic rate constants of the reactions of HOBr with pheno-late and halophenolates are on average around 3000 timesgreater than the species-specic rate constants of the analog reactions of HOCl.

Generally, a higher degree of halogenation results in lowerspecies-specic rate constants due to the electron-withdrawing effect of the halogens. However, with respectto the apparent rate constant at near neutral pH, this effect isin some cases compensated due to the lower p Ka of thehalogenated phenols ( Fig. 8). For example, the species-specicrate constant for the reaction of HOBr with 2-bromophenol(6.4 106 M 1 s 1) is more than one order of magnitudehigher than the corresponding rate constant for its reactionwith 2,6-dibromophenol (4.8 105 M 1 s 1). In contrast, theapparent second-order rate constant for the reaction of HOBrwith 2,6-dibromophenol at pH 7 (3.3 105 M 1 s 1) is higher

than the apparent rate constant for the reaction with 2-bromophenol at the same pH (2.2 105 M 1 s 1). This contra-intuitive effect is explained by the considerably lower p Ka of 2,6-dibromophenolcompared to 2-bromophenol (6.7 vs.8.5). Acomparison of the apparent second-order rate constants forthe reaction of bromine with various substituted and unsub-stituted phenols as a function of pH is shown in Fig. 8.

Quantitative Structure e Activity Relationships (QSARs) be-tween species-specic rate constants for oxidation reactionsof closely related compounds and substituent descriptor var-iables such as Hammett substituent constants ( s ) have beensuccessfully applied to predict and quantify rate constants foroxidation of organic compounds ( Canonica and Tratnyek,2003; Lee and von Gunten, 2012 ). This empirical relationshipimplies a linear correlation between the logarithm of thespecies-specic rate constant (log( k)) and the Hammett sub-stituent constant ( s ) calculated as the sum of the effects of thedifferent ring substituents. The parameter s reects the effectof the substitution on the ring on electron density by inductiveand resonance effects compared to the phenolate for which sequals zero. From the compiled data of rate constants for thereaction of HOBr with aromatic compounds ( Tables 5 and 10),a Hammett-type correlation was established (Eq. (48), Fig. 9).

log kðHOBr=PhO Þ ¼ 7:8 3:5Ss (48)

Hammett constants from Hansch et al. (1991) and Lee and vonGunten (2012) have been used. The negative slope is typical forelectrophilic reactions and its magnitude shows the sensi-tivity of the reaction to substituent effects ( Hansch et al.,1991). The correlation includes the rate constants of the re-actions of HOBr with the dissociated species of phenol, ethi-nylestradiol (EE2, Table 10), and their chlorinated andbrominated derivatives and agrees well with correlationsfound in other studies ( Acero et al., 2005; Gallard et al., 2003 ).All rate constants obtained by Guo and Lin (2009) aswell astherate constants for phenol and p-cresol from Gallard et al.(2003) have been excluded due to the relatively high uncer-tainty of these measurements (see above). Despite arriving at

Fig. 7 e Correlation of the species-specic rate constantsfor the reaction of HOBr with various phenolates with thecorresponding species-specic rate constants of HOCl.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 27



14/28

approximately the same correlation coefcients, also tri-chlorinated and tribrominated phenols were excluded fromthe correlation since the preferred sites of substitution ( ortho

and para position relative to the hydroxyl group) are notavailable and the Hammett calculation does not take thisparticular case into account.

Fig. 10 shows a comparative Hammett plot for the reactionof various oxidants including HOBr with phenolates. It showsthat only ozone reacts faster than HOBr with rate constantsthat are around one order of magnitude higher. The reactivityof HOBr is similar to the one of chlorine dioxide. While HOBrexhibits a muchhigher reactivity towardphenolic compoundsthan HOCl (ca. three orders of magnitude), the two oxidantsshow approximately the same dependence on the phenolicsubstituents moieties (slope 3.0 and 3.5 for HOCl and HOBr,respectively). HOI and to a lesser extend ferrate are more

sensitive to substitution effects, which is reected in thesteeper slopes compared to the other oxidants ( Lee et al.,2005).

Compared to phenolic compounds, aromatic compoundswithout activating or only slightly activating moieties such asmethoxyor methyl substituents (benzene,anisole andxylene)have very low rate constants for the reaction with bromine(Table 5 ).

Aromatic functional groups can make up a signicant partof NOM. The reaction of bromine with NOM typically showsbiphasic reaction kinetics: a fast initial phase followed by aslower one ( Echigo and Minear, 2006; Westerhoff et al., 2004 ).The rst phase has been attributed to the reaction with aro-

matic functional groups by means of specic UV-absorption

with apparent rate constants ranging in the order of 102e 104 M 1 s 1 (rate constants based on C ) around neutralpH, while the rate constants for the slower phase are around100 times lower ( Echigo and Minear, 2006; Westerhoff et al.,1998, 2004). However, rate constants for the reaction withNOM are difcult to compare due to the great variability of its

nature. It is unclear whether the dominant reaction mecha-nism is bromination (bromine atom transfer) or electron

Fig. 10 e Comparison of the correlations between thesecond-order rate constants of the reaction of dissociatedphenols with HFeO 4 L , O3 , ClO2 , HOI, HOCl and HOBr as afunction of the Hammett constant (sigma, s ). Reprinted(adapted) with permission from Lee et al. (2005) . Copyright (2005) American Chemical Society. Thebold line for HOBr istaken from Fig. 9.

3

4

5

6

7

6 7 8 9 10 11 12

A p p a r e n

t s e c o n

d - o r d e r r a

t e c o n s t a n

t [ M - 1 s -

1 ] , l o g s c a

l e

pH

Phenol(unsubstituted)

2-Bromophenol

4-Bromophenol

2,4-Dibromophenol

2,6-Dibromophenol

2,4,6-Tribromophenol

_ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _

Fig. 8 e Apparent second-order rate constants for thereaction of bromine with phenol, mono-, di-, andtribromophenols. At pH 7, the apparent second-order rateconstants of some of the brominated phenols are in thesame range or even higher than the corresponding rateconstant of (non-substituted) phenol, even though theirspecies-specic rate constants are more than one order of magnitude lower. This effect is due to the lower p K a of theBr-substituted compounds.

Fig. 9 e Correlation between the second-order rateconstants (log( k )) for the reaction of HOBr with dissociatedphenols and the Hammett constant (sigma, s ). Correlation:log ðk

HOBr=PhOL Þ[ 7:8L 3:5 Ss , R 2 [ 0.8. 2,4,6-tribromophenol

and 2,4,6-trichlorophenol were excluded from thecorrelation (see text).

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 228



15/28

transfer. During the initial phase of the reaction, brominationof NOM dominated with more than 75% ( Echigo and Minear,2006). In the case of preozonated NOM, however, bromina-tion occurred to 15 e 25% and electron transfer to 75 e 85% forreaction times longer than the initial phase ( Westerhoff et al.,1998).

3.3.2. Nitrogen-containing compoundsGenerally, hypohalous acids (i.e. HOCl and HOBr) reactquickly with nitrogen-containing compounds, forming varying N-halo-derivatives, which depend on the N /hypo-halous acid ratio and pH. At a ratio of 1:1 or higher, theproduct is a monohalogenated N-compound; at lower ratios,dihalogenated compounds are formed in a second reaction,which is slower than the rst halogenation ( Armesto et al.,1998).

3.3.2.1. Amines. Most of the kinetic information forHOBre amine reactions deals with primary amines; little or noinformation is available on secondary and tertiary amines(Table 6 ). The reaction of HOBr with primary and secondaryamines is very fast with apparent second-order rate constantsin the order of 10 5e 106 M 1 s 1 at near neutral pH ( Pattisonand Davies, 2004; Skaff et al., 2007; Wajon and Morris, 1982 ).A rate constant for the reaction of HOBr with glycine has beendetermined in two studies with apparent second-order rateconstants of 2.6 106 M 1 s 1 for a pH between 7.2 and 7.5(Pattison and Davies, 2004 ) and 1.4 106 M 1 s 1 for pH 7.5(Wajon and Morris, 1982 ) (Table 6 ). A difference of about afactor of two is quite common for experimentally determinedrate constants. Based on the limited data in Table 6 , it seemsthat the nature of the substituent R on the primary amine(NH2R) has no signicant effect on the magnitude of thesecond-order rate constant.

The products of the reactions of bromine with amines arebromamines formed by an electrophilic substitution of aproton by Br þ (Wajon and Morris, 1982 ). It has been found thatthe resultant organic bromamines can react further andbrominate phenolic structures ( Wu et al., 1999 ). Also the re-action of some organic bromamines with nicotinamideadenine dinucleotide (NADH) and ethylene glycol vinyl etherhas been reported: the rate constant for the reaction of brominated taurine with NADH at pH 7.3 was found to be2.7 103 M 1 s 1 and the rate constants for the reactionof brominated phosphoryl-ethanolamine, brominated phos-phoryl-serine, brominated N-a -Lysine and brominatedtaurine with ethylene glycol vinyl ether ranged between1.3 103 and w 2.7 104 M 1 s 1 at pH 7.4 (Pru ¨ tz et al., 2000;Skaff et al., 2008 ).

Due to the higher electrophilicity of HOBr compared toHOCl, the species-specic rate constants of the analogousreactions for HOBr and primary and secondary amines arearound one to two orders of magnitude higher than for HOCl(Deborde and von Gunten, 2008 ). Similar to the case of ammonia (Section 3.2.1, Fig. 6), the reaction of amines (e.g.,glycine) with bromine is not only favored due to the higherspecies-specic rate constants but also due to the differencein pKa values of HOBr (8.8) and HOCl (7.5). This difference inpKa values means that the pH range over which HOBr and theneutral form of the amine (i.e. the two reacting species) are

simultaneously present is larger than for HOCl. At the sametime, over this pH range the relative HOBr concentration isalways higher than the relative HOCl concentration.

The rate constants of the reaction of bromine with primaryand secondary amines are higher not only than the corre-sponding rate constants of chlorine, but also than the corre-sponding rate constants of ozone reactions ( von Sonntag and

von Gunten, 2012 ). The second-order rate constants for the re-actionofozonewithglycineanddimethylamineareintheorderof 105 M 1 s 1 and 10 7 M 1 s 1, respectively ( Hoigné and Bader,1983b), whereas the corresponding species-specic rate con-stants are in the order of 10 8 M 1 s 1 and 10 9 M 1 s 1 for HOBr(Table 6 ). This may affect oxidation product formation fromamines or other nitrogen-containing compounds during ozon-ation of bromide-containing waters ( von Gunten et al., 2010 ).

3.3.2.2. Other nitrogen-containing organic compounds. Similarto the corresponding reactions of HOCl ( Deborde and vonGunten, 2008 ), HOBr reacts only slowly with amides ( Table6). This can be generally expected for oxidation of amidessince the nitrogen is deactivated by the adjacent carbonylgroup ( vonSonntag and von Gunten,2012 ). In the case of urea,it was found that the chlorination was initiated by Cl 2, ratherthan HOCl (Blatchley and Cheng, 2010 ). This might also be thecase for the reaction of bromine with amides. In contrast tolinear amides, second-order rate constants for the reaction of HOBr with cyclic amides such as Cyclo(Gly) 2 or Cyclo(Ala)2have been reported to be around two orders of magnitudegreater ( Pattison and Davies, 2004; Pru ¨ tz, 1999). The bromi-nated amide Cyclo(Gly) 2 reacts further with NADH with a rateconstant of at least 4 105 M 1 s 1 (Pru ¨ tz, 1999).

Furthermore, apparent second-order rate constants for thereaction of guanidine derivatives with bromine at pH 7.2 e 7.5are in the order of 10 3 M 1 s 1. For the reaction of brominewith imidazole, indole, and pyridine moieties and urateapparent rate constants in the order of 10 6 M 1 s 1 at pH7.2e 7.5 have been reported ( Table 6 ).

3.3.3. Sulfur-containing compoundsApparent second-order rate constants for the reaction of HOBrwith sulfur-containing compounds at near neutral pH arelisted in Table 7 . In the case of the moieties of the amino acidscysteine and methionine, rate constants in the order of 107 M 1 s 1 (sulfhydryl group) and 10 6 M 1 s 1 (thioethergroup) have been reported ( Pattison and Davies, 2004; Pru ¨ tzet al., 2000). These rate constants are high, but lower thanthe corresponding apparent rate constants of the reactionwith HOCl, which are 3.0 107 M 1 s 1 for N-acetyl-cysteineand 3.8 107 M 1 s 1 for N-acetyl-methionine ( Pattison andDavies, 2001, 2004). For the reaction of HOBr with the disul-de bond in 3,3 0-dithio-dipropionic acid (DTDPA) and incystine (disulde bond), apparent second-order rate constantsin the order of 10 6 M 1 s 1 and 10 5 M 1 s 1, respectively, werereported at pH 7.2 e 7.5 (Pattison and Davies, 2004 ). The lowerrate constant for cystine compared to DTDPA might be due tosterical hindrance ( Pattison and Davies, 2004 ). The second-order rate constant for the reaction of DTDPA with HOBr isaround one order of magnitude higher than for HOCl, whilethe rate constants for cystinediffer by a factor of two ( Pattisonand Davies, 2001, 2004 ).

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 29



16/28

Table 6 e Rate constants for the reaction of nitrogen-containing compounds with bromine.

Compound Reactive Moiety p Ka Species-specicrate constants

Apparent rate constantat pH 7 or at given pH

Tem

k(HOBr)[M 1 s 1]

kðOBr ÞM 1 s 1

kapp[M 1 s 1]

N-a -Acetyl-lysine Primary amine 10.5 (3.6 0.3) 105 (pH 7.2e 7.5) Alanine Primary amine 9.9 (1.6 0.1) 106 (pH 7.2e 7.5) ε -Aminocaproic acid Primary amine (2.6 0.2) 105 (pH 7.2e 7.5) Glutamate Primary amine 10 3.5 108 5.0 104 3.0 105

Glycine Primary amine 9.8 (3.8 0.3) 108 (2.1 0.2) 105 4.6 105

(2.6 0.2) 106 (pH 7.2e 7.5) Phosphoryl-ethanolamine Primary amine (8.8 0.4) 105 (pH 7.4)

(2.6 0.4) 105 (pH 7.4) Phosphoryl-serine Primary amine (9.3 0.8) 105 (pH 7.4)

(3.1 0.4) 105 (pH 7.4) Valine Primary amine 9.7 (1.7 0.1) 106 (pH 7.2e 7.5) Dimethylamine Secondary amine 10.7 3.0 109 6.0 105

N-Acetyl-alanine Amide (7 2) 10 2 (pH 7.2e 7.5) N-Acetyl-alanine-OMe Amide 2.1 0.2 (pH 7.2e 7.5) Cyclo(Alanine) 2 Amide (2.5 0.3) 102a (pH 7.2e 7.5) Cyclo(Aspartic acid) 2 Amide (5.0 2.0) 101a (pH 7.2e 7.5) Cyclo(Glycine)2 Amide (9.0 1.0) 102a (pH 7.2e 7.5)

(2.9 0.3) 103a (pH 7.2e 7.5) 6 102a (pH 7.2)

Cyclo(Serine) 2 Amide (5.5 0.9) 102a (pH 7.2e 7.5) 2-Methylpropionamide Amide 1.5 0.3 (pH 7.2e 7.5) Propionamide Amide 3.3 0.4 (pH 7.2e 7.5) Trimethylacetamide Amide (9 1) 10 1 (pH 7.2e 7.5) N,N-Dimethylsulfamide Sulfamide 10.5 8.1 108 2.7 105

Urate Guanine (3.4 0.2) 106 (pH 7.4) N-Acetyl-arginine-OMe Guanidine (2.2 0.4) 103 (pH 7.2e 7.5) Ethylguanidine Guanidine (1.3 0.2) 103 (pH 7.2e 7.5) N-Acetyl-tryptophan Indole (3.7 0.3) 106 (pH 7.2e 7.5)

4-Imidazoleacetic acid Imidazole w

3.0 106

(pH 7.2e

7.5) NMNH (Nicotinamide mononucleotide) Pyridine w 4 106 (pH 7.2) NMNH, brominated Br-Pyridine w 9 105 (pH 7.5)

a Rate constant given per amide group, i.e. the rate constant for the molecule would double ( Pattison and Davies, 2004 ).

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


17/28

3.3.4. OlenesTable 8 provides the available kinetic information for the re-action of olenes with HOBr. At pH 7.4, ascorbate has a highapparent second-order rate constant because one hydroxylicgroup is deprotonated, which signicantly increases theelectron density in the double bond. The double bond in

ethylene glycol vinyl ether reacts very rapidly with HOBr,while the corresponding rate constants for the reaction with3-pentenoic acid and sorbate are about two and three ordersof magnitude lower ( Table 8 ). The lower rate constant for thereaction with sorbate compared to 3-pentenoic acid can beexplained by a conjugation of the electron-withdrawing carboxyl group in the case of sorbate.

3.3.5. Carboxylic acids, aldehydes, alcohols, and ketonesThe reported apparent second-order rate constants for thereaction of HOBr with carboxylic acids are low. Formic andacetic acid are practically unreactive, while the dicarboxylic

acids oxalic and malonic acid show a low reactivity ( Table 9 ).Even lower rate constants have been reported for aldehyde(formaldehyde, pH 4) and alcohol (2-propanol, pH 6.7)(Table 9 ). Rate constants for the reaction of Br 2 with acetal-dehyde and the alcohols 2-propanol and ethanol near neutralpH have been reported to be in the order of 10 1 M 1 s 1

and 104

M1

s1

, respectively ( Perlmutter-Hayman andWeissmann, 1962, 1969 ). Species-specic rate constants forthe reaction of OBr for various alcohols range between4.1 10 7 and 3.4 10 4 M 1 s 1 (Negi et al., 1987).

Ketones only react with HOBr in their enol form ( Fig. 11).For the reaction of HOBr with ketones, their enolization is therate-determining step, because the reaction of the enol (ole-ne, see above) with bromine is fast ( Pinkernell and vonGunten, 2001 ). The ketone-enol equilibrium is generallyheavily on the ketone side and enolization rate constants arelow; in the order of 10 6 s 1 for aliphatic ketones in acidicmedia ( Dubois and Toullec, 1969 ). Similar to the bromination

Table 9 e Rate constants for the reaction of carboxylic acids, aldehydes, and alcohols with bromine.

Compound p Ka Apparent rate constant Temperature [ C] Reference

kapp [M 1 s 1]Acetic acid 4.8 < 1 (pH 6) 20 (Pinkernell and von Gunten, 2001 )Formic acid 3.8 < 1 (pH 6) 20 (Pinkernell and von Gunten, 2001 )Malonic acid 2.8 3.0 101 (pH 4) 20 (Pinkernell and von Gunten, 2001 )

5.7Oxalic acid 1.3 4.0 101 (pH 6) 20 (Pinkernell and von Gunten, 2001 )

4.1Formaldehyde 1.8 10 3 (pH 4) 20 (Pinkernell and von Gunten, 2001 )2-Propanol 3.9 10 4 (pH 6.7) 25 (Perlmutter-Hayman and Weissmann, 1969 )

Table 7 e Rate constants for the reaction of sulfur-containing compounds with bromine.

Compound p Ka Apparent rate constantat pH 7.2e 7.5

Temperature [ C] Reference

kapp [M 1 s 1]

N-Acetyl-cysteine 8.1 ( e SH) (1.2 0.2) 107 22 (Pattison and Davies, 2004 )N-Acetyl-cystine ( N-Acetyl-Cysteine) 2 (3.4 0.8) 105 22 (Pattison and Davies, 2004 )N-Acetyl-methionine-OMe (3.6 0.3) 106 22 (Pattison and Davies, 2004 )

(9.6 1.3) 106 37 (Pattison and Davies, 2004 )3,30-Dithio-dipropionic acid (DTDPA) (1.1 0.2) 106 22 (Pattison and Davies, 2004 )

(2.3 0.2) 106 37 (Pattison and Davies, 2004 )Methionine w 4 106 (pH 7.2) 20 2 (Pru ¨ tz et al., 2000 )

Table 8 e Rate constants for the reaction of olenes with bromine.

Compound p Ka Apparent rate constant at pH 7.4 Temperature [ C] Reference

kapp [M 1 s 1]Ascorbate 4.1 (1.7 0.2) 106 22 (Skaff et al., 2007 )

11.8Ethylene glycol

vinyl ether(3.5 0.1) 106 22 (Skaff et al., 2008 )

3-pentenoic acid 4.5 (1.1 0.2) 104 22 (Skaff et al., 2007 )Sorbate 4.8 (1.3 0.2) 103 22 (Skaff et al., 2007 )

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 2 31



18/28

of aromatic compounds, bromination of enols (olenes) is animportant source of THMs.

3.3.6. MicropollutantsThe reactivity of HOBr with several micropollutants has beeninvestigated and depends on the functional groups present inthe micropollutant ( Table 10 ). Fig. 12 shows the chemicalstructures of micropollutants for which data on their reac-tivity with bromine is available, including the likely site of reaction with bromine.

3.3.6.1. Aromatic compounds. 17a -Ethinylestradiol (EE2) andits halogenated transformation products have high second-order rate constants for their reactions with HOBr, due totheir phenolic moieties ( Table 10 ) (Lee and von Gunten, 2009 ).The species-specic rate constants decrease with increasing degree of halogenation of the phenolic ring (for discussion seeSection 3.3.1). Since the reactivity of HOBr is about three or-ders of magnitude higher than the one of HOCl (see Fig. 7), thetransformation of EE2 is accelerated during chlorination of bromide-containing waters ( Lee and von Gunten, 2009 ). It hasbeen shown that bromination of EE2 is an effective means todestroy the estrogenic properties of EE2 ( Lee et al., 2008).Endpoints other than estrogenicity have not yet been inves-tigated for brominated EE2. Chlorophene is as well attacked atits phenolic moiety. Its apparent rate constant at pH 7(1.9 102 M 1 s 1) is a factor 100 lower than what could beexpected by comparison to the rate constant of 4-chlorophenol (3.2 104 M 1 s 1, Table 5 ).

In contrast, the phenylurea herbicides diuron and iso-proturon react only slowly with bromine ( Acero et al., 2007 )(Table 10 ), because the aromatic ring is deactivated (diuron) oronly slightly activated (isoproturon) and nitrogen is present asan amide, which is characterized by a low reactivity withHOBr (see Section 3.3.2). DEET shows an apparent rate con-stant comparable to diuron, which is explained by its deacti-vated aromatic ring. Phenacetin is an aromatic compoundwith only a slight activation by an ethoxy group, which resultsin an apparent second-order rate constant of 7.3 M 1 s 1 for itsreaction with bromine at pH 7 ( Table 10 ). Naproxen has ahigher reactivity due to the presence of a naphthalene moiety(k ¼ 25 M 1 s 1, Table 10 ). This is comparable to the case of ozone, whose second-order rate constant forthe reaction with

naphthalene (3 103 M 1 s 1) is 1500 times greater than therate constant for its reaction with benzene (2 M 1 s 1) (Hoignéand Bader, 1983a; von Sonntag and von Gunten, 2012 ).

3.3.6.2. Nitrogen-containing compounds. The reaction sites forpotential bromine attack on metoprolol are the aromatic ring or the secondary amine group. Based on the reactivity of

dimethylamine (apparent second-order rate constant of around 10 5 M 1 s 1 at pH 7, Table 6 ) it would be expected thatthe secondary amine moiety dominates the kinetics. How-ever, there is a discrepancy between the low rate constant forthe reaction with metoprolol at pH 7 ( Table 10 ) and the re-ported rate constant for the reaction with dimethylamine(Table 6 ). The same is true for nortriptyline and hydrochlo-rothiazide (its aromatic ring is unreactive to electrophilicattack due to the electron-withdrawing chloro- andsulfoxide-moieties), where the attack is also expected at thesecondary amine group. More data is needed to interpretthese ndings.

As far as attacks on heterocyclic nitrogen are concerned,bromine shows a very high reactivity with 3-methylindole(1.1 108 M 1 s 1 at pH 7). A high rate constant was alsofound for the reaction of the indole moiety in the amino acidtryptophan (3.7 106 M 1 s 1 at pH 7.2e 7.5, Table 6 ). Thereactivity of bromine with benzotriazole is rather low(8.5 M 1 s 1 at pH 7).

3.3.6.3. Sulfur-containing compounds. The pesticide chlor-pyrifos shows a high second-order rate constant for the re-action with bromine ( Table 10 ), which is three orders of magnitude greater than the corresponding rate constant of the reaction with chlorine (4.8 102 M 1 s 1) (Duirk et al.,2008). Because the aromatic ring is highly deactivated bychlorine substitution, the thiophosphate group is oxidized,forming chlorpyrifos oxon, which is more toxic than theparent compound ( Wu and Laird, 2003 ).It was found thatOBrand OCl accelerate the hydrolysis of both the parent com-pound and the oxon to form 3,5,6-trichloro-2-pyridinol ( Duirkand Collette, 2006; Duirk et al., 2008 ). The rate constants forthe hydrolysis by OBr and OCl range between 0.27 and0.39 M 1 s 1 (Duirk and Collette, 2006; Duirk et al., 2008 ).

The species-specic rate constants for the reaction of bromine and chlorine with ametryn were reported to behigher for the positively charged, protonated species than forthe neutral species of ametryn ( Xu et al., 2009). However, thereason for an increasing apparent rate constant withdecreasing pH is probably not the speciation of ametryn andthus not a higher species-specic rate constant for the reac-tion with protonated ametryn compared to the neutral form.This effect is more likely due to the contribution of Br 2 at lowpH. As observed for the other compounds as well, the second-order rate constant of the reaction of bromine with ametrynwas signicantly higher than that for chlorine ( Xu et al., 2009).An attack on the sulfur, resulting in a sulfoxide was proposedfor the reaction of chlorine ( Xu et al., 2009); the same mech-anism can be assumed for bromine, because the rest of themolecule (triazine ring and deactivated secondary amines) isnot susceptible to oxidative attack ( von Sonntag and vonGunten, 2012 ). The apparent second-order rate constant forthe reaction of bromine with ametryn at pH 7 is more than two

C C

O

R3R2

R1

H

C C

OH

R3

R1

R2

Ketone Enol

HOBr

Fig. 11 e Enolization of ketones followed by reaction with bromine. Bromine attack occurs at the double bond of theenol (arrow). The rate-determining step of this reaction isthe enolization.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 232



19/28

Table 10 e Rate constants for the reaction of micropollutants with bromine.

Compound p Ka Species-specic rate constants Apparent rateconstant at pH 7

Te

kðHOBrþ Hþ Þa kðHOBr=Pþ Þb k(HOBr/P) b kðHOBr=P Þb kðHOBr=P2 Þ

b kðOBr =P Þb kapp[M 2 s 1] [M 1 s 1] [M 1 s 1] [M 1 s 1] [M 1 s 1] [M 1 s 1] [M 1 s 1]

Ametryn 4.1 3.5 106 9.1 103 1. 3 104

Amoxicillin 2.6 < 1 107 (2.4 0.4) 104 (2.4 0.4) 104 (9.9 1.5) 106 (3.8 1.2) 109 6.7 106

7.39.7

Benzotriazole 0.4 8.5 0.3 8.2

Chlorophene 9.8 (1.9 0.2) 102

Chlorpyrifos (3.2 0.6) 105 3.1 105

DEET 0.67 (9 2) 10 2

Diuron 1.2 105 1.2 10 1 8 10 2 1.3 10 1

EE2 10.4 (7.7 2.2) 103 (5.3 0.7) 108 (1.2 0.8) 106 2.1 105

2-Cl EE2 8.7 (7.2 1.3) 107

1.5 106

4-Cl EE2 8.6 (2.8 1.4) 107 6.6 105

2-Br EE2 8.9 (1.3 0.4) 108 1.7 106

4-Br EE2 9.0 (7.4 2.6) 107 7.0 105

2,4-DiCl EE2 7.1 (4.9 1.2) 105 2.1 105

2,4-Br,Cl EE2 7.1 7.3 105 3.2 105

2,4-Cl,Br EE2 7.1 7.3 105 3.2 105

2,4-DiBr EE2 7.1 (9.7 2.4) 105 4.1 105

Hydrochlorothiazide 7.9 (4.2 0.8) 103 3.3 0.4 (5.8 1.0) 101 8.1 1.2 Isoproturon 3.1 107 1.8 101 2 10 2 2.5 101

Metoprolol 9.7 (2.8 0.2) 103 3.9 0.2 (1.0 0.1) 102 5.2 0.8 3-Methylindole (1.1 0.1) 108

Naproxen 4.2 < 1 108 (2.1 0.4) 104 (2.7 0.3) 102 (2.5 0.1) 101

Nortriptyline 10.2 7.4 0.3 Phenacetin 1.4 < 1 104 (5.4 0.6) 102 1.7 0.2 (7.4 0.5) 107 7.3 1.0

14a It is not clear, whether these rate constants truly characterize an acid-catalysis or rather the effect of other bromine species formed at low pH (see Sectionb Pþ , P, P and P2 represent the charge of the molecule and not in each case its actual protonation.

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


20/28

Fig. 12 e Structures of micropollutants for which second-order rate constants for the reaction with bromine are available (in

alphabetical order). Arrows indicate the possible reaction sites for bromine.

w a t e r r e s e a r c h 4 8 ( 2 0 1 4 ) 1 5 e 4 234



21/28

orders of magnitude smaller than the corresponding rateconstant for N-acetyl-methionine-OMe with a similar reactivemoiety (thioether, Table 7 ). The lower reactivity of ametryncan be explained by the electron-withdrawing effect of thetriazine moiety.

For amoxicillin, an attack of bromine is possible at thephenolic, the amine or the sulfur moiety. Based on the reac-

tivity of the different functional groups, the phenolic moietycan be ruled out. It also seems that the second-order rateconstants for the reaction of HOBr with the primary amine(see Table 6 ) are too low to explain the apparent rate constantof the reaction of bromine with amoxicillin at pH 7. Therefore,it is hypothesized that the primary attackof bromineoccursatthe thioether group.

4. Modeling of chlorination processes in thepresence of bromide

To assess the effect of bromide on oxidative transformationreactions during chlorination, a kinetic model was set up,which is illustrated in Fig. 13. It combines the formation of HOBr through the oxidation of bromide by HOCl (Eq.

Documents

Oxidative treatment of bromide containing waters formation of bromine and its reactios with inorganic and organic.pdf