Environ. Sci. Techno/. 1988, 22, 1049-1056
Reactivities of Hypochlorous and Hypobromous Acid, Chlorine Monoxide, Hypobromous Acidium Ion, Chlorine, Bromine, and Bromine Chloride in Electrophilic Aromatic Substitution Reactions with p -Xylene in Water Evangelos A. Voudriast and Martin Reinhard" Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, California 94305
Scheme I
w Rate laws for halogenation of electrophilic p-xylene (RH) to RX (X = C1 or Br) were determined in dilute solutions of aqueous chlorine and bromine at 20 "C in the pH range 2.5-7.5 and at ionic strength 0.1 M. For the electrophiles (X-Y), BrC1, C12, C120,Br2, and HOBr, disappearance of RH was found to follow the second-order rate equation -d[RH]/dt = kq,-y)[RH] [X-Yl
Scheme I1
The specific rate constants, k2, decreased in the following order: BrCl (540 f 9 M-l s-l), C12 (58 f 9 M-l s-l), C120 (54 f 4 M-l s-l), Br, (0.251 f 0.004 M-l s-l), and HOBr [(1.1f 0.7) X M-l s-l]. The k2 value of HOCl was approximately 0 and relatively insignificant. A third-order term -d[RHI /dt = ~ ~ ( H ~ O B ~[HOBr] + ) [ H +[RH] I was identified for the acid-catalyzed bromination with HOBr with k3(H,0Br+)= 1330 f 10 M-2 .'-S Introduction Hypohalous acids HOX (X = C1 or Br) are relatively weak electrophiles which react only with aromatic substrates having activated rings (1,2). However, protons and some anions, such as OC1-, C1-, or Br-, may convert HOX into highly reactive species and, thus, may catalyze halogenations of relatively nonreactive substrates (1-6). The reactive species X-Y may be a permanent dipole consisting of partially positively charged X6+attached to a partially negatively charged carrier Yh. X may be Br or C1, and Y may be H20, HO, Br, Cl, or OC1. In homoatomic species (Cl,, Br2) X6+-X6-is formed by dipole induction (1). In acidic aqueous HOCl solutions, possible chlorinating agents are HOC1, Cl2O,H20C1+(1,6),and, if C1-is present, Clz (2, 7,8). In acidic aqueous HOBr solutions possible brominating agents are HOBr and H20Br+and, if Br- or C1- is present, Br, and BrC1, respectively (1,3,5). Chlorinating and brominating agents may coexist in mixed systems, which contain HOCl and Br-. In such system bromination and chlorination may be competing reactions as will be shown in the following paper (9). These systems are of public health interest because chlorine disinfection of water containing organic compounds and Br- may lead to the formation of chlorinated and brominated aromatic byproducts (7, 9, 10). The mechanism of aromatic electrophilic substitution reactions is thought to proceed through Wheland intermediates (also termed CT complexes or arenium ions) according to Scheme I (11). Formation of a Wheland intermediate is generally preceded by rapid bimolecular formation of an encounter or a a-complex [X-Y, RH], where RH is the aromatic compound. The encounter complex of phenol and C12 [Clz,CGH50H],for example, is formed at a rate of (2.3 f 0.5) X lo4 M-l s-l (8). ~
~~
~~~
'Present address: School of Civil Engineering, Georgia Institute of Technology, Atlanta, GA 30332. 0013-936X/88/0922-1049$01.50/0
If formation of the Wheland intermediates is rate-limiting, the rate equation for the disappearance of RH is (11) -d[RH]/dt = k,[RH][X-Y]
(1)
If reactive X-Y is formed in a rapid preequilibrium from the nonreactive HOX according to HOX
+ Y- + H+ z% kh
X-Y
+ H20
(2)
the observed rate due to reaction with X-Y is of the fourth order (2): -d[RH] /dt
k4[H+][HOX] [Y-][RH]
(3)
where [HOX] is the instantaneous concentration of HOX. If the reaction is acid-catalyzed and Y- ions are absent, the observed overall rate is of third order (3-6): -d[RH]/dt = k,[HOX][H+][RH]
(4)
Gilow and Ridd (3)have formulated Scheme I1 to interpret this third-order term. It considers two mechanistic alternatives for the formation of [RH, H20Br+] either by preformation of H20Br+or, alternately, by protonation of the encounter or a-complex [RH, HOBr]. They speculated that preformation of H20Br+is relatively dominant for nonactivated substrates at high acidities and less dominant for activated substrates at low acidities. In both cases, H20Br+may be considered as the electrophilic species, although the exact mode of HzOBr+ formation is not known. The objective of this study was (i) to determine the rate laws of aromatic p-xylene halogenation and (ii) to determine the reactivities of all possible electrophiles present in aqueous bromine and chlorine solutions. To date, limited information is available for a quantitative comparison of electrophile reactivities. p-Xylene was used as the substrate because (i) it forms only one product, 2-X-pxylene, upon monohalogenation [which circumvents the problem of changing product distribution with changing electrophile selectivity (S)],(ii) compared to benzene and other methylated benzenes, p-xylene exhibits an intermediate reactivity, and (iii) substitution is possible only
0 1988 American Chemical Society
Environ. Sci. Technol., Vol. 22, No. 9, 1988
1049
in ortho position to methyl groups which have nu adjacent methyl groups (“single methyl groups”). Such single methyl groups exhibit little steric hindrance. The relative chlorination rates of benzene:p-xy1ene:pentamethylbenzene are 5 X 10“:1:3.6 X lo4 (12). Therefore, a relatively broad range of electrophile reactivity can be compared directly by using a single substrate. Alternate approaches employ substrates having a range of reactivities and development of the Hammet constant or the selectivity relationship (11, 13). Steric hindrance between substituents and the electrophilic species may’affect the reaction rates if the electrophiles have large steric requirements and if substitution occurs in the position ortho to “double-ortho” or “buttressed single-ortho” methyl groups. However, pxylene is substituted in a position ortho to a single methyl group (“single-ortho”),where such effects have been shown to be small (14). Water was selected as the solvent because aqueous rate and equilibrium constants of many of the reactions which convert HOX into electrophiles of greater reactivity are known. Br2 and C12may be formed according to eq 2 with X = Y = Br and C1, respectively. Mixed systems containing C12, Br2 and the mixed halogen BrCl may be formed when HOCl or C12is added to water containing Br-. First, Br- is rapidly and irreversibly oxidized t o HOBr:
the protonation of HOBr and HOCl (eq 12 and 13, respectively). Experimental Section
Materials. The water used in all experiments was prepared from deionized water further purified with a MilliQ system (Millipore Corp., Bedford, MA). HOCl and HOBr concentrations were determined by the DPD titrimetric procedure (25). Aqueous solutions of hypobromous acid were prepared by adding liquid bromine (J. T. Baker Chemical Co., Phillipsburg, NJ) to water containing an excess of freshly prepared silver phosphate (5), the latter having been produced by mixing stoichiometric amounts of AgN03 and Na2HP04. The suspension (-350 mL) was shaken at intervals of 30 min. After approximately 2 h, HOBr was distilled at 38-42 “C under vacuum from a water aspirator. The first fraction (80-100 mL) contained a high Br- concentration and was discarded. The second Br--free fraction (120-150 mL) was stored at 4 “C in the dark until used. The concentration of the second fraction was 0.015-0.02 M and did not show a precipitate upon addition of Ag+. The diluted solutions were stable in the course of 5-6 h as measured by titrimetry. Aqueous solutions of hypochlorous acid were prepared by bubbling ultrahigh-purity chlorine gas (Matheson, HOCl Br- A HOBr C1(5) Newark, CA) into water containing mercuric oxide (HgO) and freshly precipitated silver phosphate (6). After apwith k = 2.95 X lo3M-l s-l at 25 “C (15). Higher oxidation proximately 2 h of occasional shaking, aqueous HOCl products of Br- are possible byproducts. At pH >10 the (similar volumes as HOBr) was distilled at 28-30 “C under predominant product is OBr- with small amounts of brothe vacuum of an aspirator. The first fraction contained mate (Br03-) and chlorate (C103-) (16). In the pH range a high C1- concentration as indicated by AgCl precipitation 7-10, mixtures of hypochlorite and hypobromite are and was discarded. The second HOCl fraction (0.04-0.05 unstable with respect to chlorate and bromate (16). Under M) was tested with AgN03 and was found to contain traces the conditions of this study (low concentrations in the of C1-. Diluted aliquots of this practically “Cl--free” dark), the halate formation is slow and may be ignored fraction (Cl- concentration lO-*-lO-’ M) were stored at 4 (17). “C in the dark until used. When C1- is added to an acidic HOBr solution, the 2-Bromo-p-xylene (>98.6% pure) was prepared by remixed halogen bromine chloride (BrC1) is formed (5,18). acting 0.5 mol of p-xylene with 0.5 mol of Br2in CC14 (26). The formation constant Kf(BrC]) is given by It was purified from the reaction mixture by fractional HOBr H+ C1- @ BrCl + H 2 0 distillation and characterized by gas chromatography (GC) (6) and gas chromatography/mass spectrometry (GC/MS). Kf(BrC1) = [BrClI/[HOBrI [ H + l r ~ W - l r ~ l Kinetic Experiments. The aqueous halogen solution was placed in 60-mL Wheaton bottles with Teflon seals, Because only a 0 “C value was found in the literature (18), leaving approximately 0.5 mL headspace to facilitate we determined &(BrC1) at 20 “c(19) using the electroconmixing. The ionic strength of the reaction solutions was ductivity (EC) method previously employed for Br2 (20) adjusted to 0.1 M with NaNO,. The reaction was typically and C12 (21). The chemical and thermodynamic properties started by injecting 10 pL of a 0.2044 M p-xylene solution of BrCl resemble those of C12 and Br2. The formation in p-dioxane to give an initial concentration of 3.4 X constant of BrCl is intermediate but closer to that of Clt. >>~K ~ ( B > ~ c ~ ) M. In most experiments, no buffer was used and no The Kf values decrease in the order K ~ () B measurable pH change before and after the experiment Kf(C12) with the K values of BrCl and C1, dijfering by only was observed. In some case, a pH change of 0.01-0.02 pH a factor of 11. unit was observed, in which case the average [H+] conThe equilibria which have been considered in this study centration was used for calculations. A total of 1 X are summarized in Table I. They include the formation M phosphate buffer had no apparent effect on the rate of BrC1, C12, and Br2 (eq 6, 7, and 8, respectively); the constants. After a thorough mixing, the bottles were ionization of HOCl and HOBr (eq 9 and 10, respectively); placed in a 20 f 0.01 “C water bath in the dark. The the dehydration of HOCl to chlorine monoxide (eq 11); and
+
+
+
+
Table I. Formation and Dissociation Constants of Chlorinating and Brominating Agents in Water
1050
reaction
constant
HOBr + H+ + C1- BrCl + Hz0 HOCl + H+ + C1- zClz + H20 HOBr + H+ + Br- 3 Brz + HZ0 HOCl e? H+ + OClHOBr e? H+ + OBrHOCl + HOCl F! ClzO + HzO HOBr + H+ F! HzOBr+ HOCl + H+ s H20Cl+
Kf(BrC1) = 3.16 X lo4 M-’ (20 OC) K ~ (= ~2.93~ x~ 103 ) M-2 (20 oc) Kf(BrZ) = 2.53 X IO8 M-’ (20 “C) pK, = 7.58 (20 OC) pK, = 8.70 (25 OC) K ~ (= 3.57 ~ ~x ~10-3~M-1, (20 o c ) Kf(Hz~~r+) N lo-’’ M-’ K~(H~OC~+) N M-’
Envlron. Sci. Technol., Vol. 22, No. 9, 1988
ref 19 21
20 22
20 19, 23 24 1
eq no. 6 7
8 9 10 11 12 13
Table 11. Typical Mass Balance Data for the Bromination or Chlorination of p-Xylene in Aqueous HOBr or HOCl Solutions reaction conditions'
reaction time, I
M, pH 6.53, [Br-I, = 0 [HOBr], = 0.124 X same as above same as above M, pH 2.95, [Br-I, = 0 [HOBr], = 0.221 X M, pH 2.95, [Br-I, = 0 [HOBr], = 0.320 X M [HOBr], = 0.207 X lom3M, pH 5.95, [Br-1, = 1 X same as above M same as above but [Br-1, = 1 X same as above same as above but [Br-I, = 5 X 10"' M M, pH 3.65, [Cl-1, = 1 X M [HOBr], = 0.223 X same as above but [Cl-1, = 1 x M [HOBr], = 0.801 x M, pH 3.60, [Cl-1, = 1 X IO" M same as above but [Cl-1, = 5 X 10"' M [HOCl], = 0.451 X lou3M, pH 6.08, [Cl-1, = 1 X lo4 M M M, pH 6.08, [Cl-1, = 1 X [HOCl], = 0.451 X
6 180 8 400 9 600 4 200 2 220 6 360 13500 18480 5 760 5 880 6 480 1260 1200 1500 15540 56 940
initial mol 1.87 X 1.87 x 1.87 x 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X 1.87 X
remaining RH, mol
lo* 10" 10-6 lo4 lo* 10" lo* lo* IO* lo*
lo*
lo* lo4 lo* 10" lo*
1.85 X 1.89 X 1.90 x 3.13 X 5.49 X 1.91 X 1.72 X 1.40 X 1.71 X 1.71 X 6.70 X 6.83 X 6.99 X 1.35 X 1.96 X 1.91 X
ndc nd 10-6 3.74 x 1.49 X lo-' lo-' 1.32 X 10" 4.01 X 7.05 X 10" 4.08 X lo* lo* 1.44 X lo* 8.67 X 1.29 X 1.20 X lo-' 1.80 X 1.74 X lo-' 10" nd 5.58 X lo*
[RH], = [RH]i + [RX]i
(14)
where [RH], = initial p-xylene concentration, [RHIi = unreacted p-xylene concentration at time t, and [RXIi = sum of mono- and dihalogenated p-xylene concentration at time t. The subscript i indicates an individual data
[ (RH
X
+ RX)/RH,] x 100 98.9 101.1 101.8 96.4 99.9 104.3 95.7 96.7 99.1 96.1 104.8 100.7 99.9 100.3 104.8 102.4
10" lo*
'All experiments were conducted at 20 OC in the dark in 60-mL Wheaton bottles; [RH], = 0.34 for the 5 mL of reaction solution withdrawn at the end of exDeriment. nd = not detected.
reaction was stopped by injecting 5 mL of 0.1 N Na2S03 solution (while simultaneously withdrawing 5 mL of reaction solution). The reaction mixtures were extracted into 2 mL of hexane and analyzed by gas chromatography (Carlo-Erba, 30-m DB-5 fused silica capillary column, flame-ionization detection, temperature program 50-170 "C at 5 OC/min). The mechanism of p-xylene halogenation was studied by the method of isolation. Product formation was determined under conditions of pseudo-first-order kinetics with respect to HOCl or HOBr as a function of [H'], [Cl-1, and [Br-1. Conditions were selected such that either chlorinated or brominated products, but no mixed products, were formed. The reactions were stopped before significant amounts of dichloro- and dibromoxyleneswere formed. In case small amounts of these products were present, their peak area was added to the peak area of monohalogenated xylenes, and the same GC response factor on molar basis was assumed. Light, which promotes side-chain halogenation, was excluded from the reaction bottles except during filling of the bottles. No benzyl halides or their hydrolysis products were detected. Brown and Stock (12)reported less than 1-2% aromatic methyl group chlorination in acetic acid, and therefore, side-chain chlorination was deemed even less significant in water which is more polar. In all experiments conducted at pH > [RH]. kl is proportional to [HOBr], and [H'I, (Figures 2 and 3, respectively), which suggests a rate equation of the form -d[RHl/dt = k 3 ( ~ z ~ ~ r +[HOBrl ) [ H f l [RHI
(20)
The rate constant k3(HzOBr+)obtained by linear regression has the value of 1330 f 10 M-2 s-l. Equation 20 is consistent with the mechanisms indicated in Scheme I1 (3). Bromination in the System HOBr/H+/Cl-. The catalytic effects of Cl- and H+ were evaluated by measuring the reaction rates under pseudo-first-order conditions with respect to [RH] as a function of [HOBr], [H+], and [Cl-1. 1052 Environ. Sci. Technol., Vol. 22, No. 9, 1988
0
~
02
04
~
0.6
08
I0
~
~
[ H O B r I o , M x103
Figure 5. Dependence of slope of Figure 4, SI(,,,, r2
= 0.999.
upon [HOBr],,
The data summarized in Figure 4 suggest a two-term rate law: kT = k3(HzOBr+)[H+1 [HOBr] + Sl(BrCI)[C1-1
(21)
where kT is the overall pseudo-first-order rate constant, k 3 (OB~+)[H+][HOB~] ~ is the term due to bromination by H20br+,and Sl(BS\)[cl-]is a term first order in Cl-. ~ ~ ( B s I ) is also first order in [HOBr] as indicated in Figure 5. The overall pseudo-first-orderterm kT is first-order with respect
I
I
I
I
I
1
I
1
1
1
1
IO-
[H'],
M x103
Flgure 6. Dependence of pseudo-first-order rate constant k , for the chloride-catalyzed bromination of p -xylene in aqueous HOBr solution upon [H']: [HOBr], = 0.161 X lom3M, [Cl-1, = 5 X I O 4 M, r 2 = 0.995. Error bars represent 95% confidence interval (not shown if smaller than size of symbol).
2e
00
2
4
6
8
1
0
[Br-Id M x IO4
Flgure 7. Dependence of pseudo-first-order rate constant k , for the bromlde-catalyzed bromination of p -xylene in aqueous HOBr solution, pH 5.8, upon [Br-1,. (A) [HOBr], = 1.09 X M, r 2 = 0.998. (B) [HOBr], = 0.583 X lo3 M, r 2 = 0.998. (C) [HOBr], = 0.209 X lo3 M, r 2 = 0.998. Error bars represent 95% confidence interval (not shown if smaller than size of symbol).
where the fourth-order term incorporates the equilibrium equation for BrCl (eq 6) and is consistent with involvement of BrCl as the brominating agent. The observed rate constant was evaluated staiistically as k4(BrC1)= (1.08 f 0.02) X lo7 M-3 s-l. Bromination in the System HOBr/H+/Br-. The catalytic effect of Br- on bromination was evaluated under pseudo-first-order conditions with respect to [RH] as a function of [HOBr], [H+], and [Br-1. Data evaluation followed the same approach as with the system HOBr/ H+/Cl-. The data summarized in Figure 7 show the firsborder dependence of kT upon [Br-1, at three different HOBr concentrations. The first-order dependence on [HOBr], is indicated by the data shown in Figure 8 which depicts the slope of Figure 7 as a function of [HOBr],. In analogy to the HOBr/H+/Cl- system, first-order dependence on [H+]was assumed, and the following rate law was proposed:
+
-d[RH]/dt = K3(H,,B,+)[H+][HOBr][RH] ~ ~ ( B , ) [ H O[H+l B ~ [Br-] ] [RH] (23) where the fourth-order term is consistent with the involvement of Br2as the brominating agent. The observed
[Br-lo,
4
M x IO
Flgure 9. Dependence of pseudo-first-order rate constant k , for the bromide-catalyzed bromination of p -xylene in aqueous HOBr upon [Br-1,: [HOBr], = 0.891 X M, pH 7.2,r 2 = 0.998. Error bars represent 95% confidence interval (not shown is smaller than size of symbol).
rate constant was evaluated statistically as k4(Br2) = (3.98 f 0.01) X lo7 M-3 s-l. To examine the term kz(HoB,)[HOBr][RH] due to bromination by HOBr, pH conditions were selected where H+ catalysis is relatively less significant. Figure 9 shows kT at pH 7.2 and [HOBr], = 0.89 X lo9 M at increasing [Br-1, concentrations. These data were evaluated by applying the rate law -d[RH]/dt = ~~(H,oB,+)[HOB~][H+][RH] + k4(Br2)[HoBrl LH+l [Br-l [RH1 + k2(HOBr)[HOBrl[RH] (24)
Br- was included in the evaluation in order to obtain reaction rates sufficiently fast for analysis and to alleviate problems with oxidation as a side reaction. Added bromide also removed uncertainties stemming from residual effects of Br-. The constant k 2 HOBr) was evaluated statistically as (1.1f 0.7) X M-j s-l. Chlorination in the System HOCl/H+/Cl-. The catalytic effect of C1- on chlorination was determined by measuring k~ as a function of [Cl-1. The data depicted in Figure 10 indicate a first-order term in [Cl-] as was found previously (7). Since the chlorination rate at the extrapolated intercept [Cl-1 = 0 is significant and depends linearly on [HOC1I2 (Figure ll), a rate law that includes chlorination by C12and, in parallel, by C120can be considered: -d[RH]/dt = ~ ~ c I , ) [ H O[H' C I~ IfC1-1 [RH]+ kqci20)[HOC112[RHl(25) The statistical evaluation of the data using nonlinear multivariate regression analysis yielded a k4(C12) of (1.06 f 0.17) X lo6 M-3 5-l and a k3(C120) of 0.19 f 0.01 M-2 s-l. Environ. Sci. Technol., Vol. 22, No. 9, 1988
1053
5 C '
5t
"
i i i
'
-I
I
K
J .
O L Y
h
$
C
b 6 Q IO' [CI-I,, ~ ~ 1
0
0I f o 1 0
,
l o , 4
2
I
[HOCI]
~
Flgure 10. Dependence of pseudo-first-order rate constant k , for the chloride-catalyzed chlorination of p-xylene in aqueous HOC1 solution, pH 6.1, upon [Cl-1,. (A) [HOCI], = 1.03 X M, r 2 = 0.984. (B) [HOCI], = 0.728 X lov3M, r 2 = 0.998. (C) [HOCI], = 0.451 X M, r 2 = 0.997. Error bars represent 95% confidence interval (not shown if smaller than size of symbols.
, , 6
, 8
IO
,! 12
,M2x 10'
Flgure 12. Dependence of pseudo-first-order rate constant k , for the chlorination of p-xylene in aqueous HOC1 upon the square of [HOCI],. [CI-] = 0. Error bars represent 95% confidence interval (not shown if smaller than size of symbol).
The bromination and chlorination of p-xylene with HOBr and HOC1, respectively, were found to be slow or insignificant relative to the reactions catalyzed by H+, Cl-, or Br-. Acid-catalyzed brominations in the absence of Xwere observed to be first-order in HOBr, H+, and RH (eq 20), which is consistent with both mechanisms shown in Scheme 11. The significance of preformed H20Br+as the effective electrophile may be evaluated considering eq 12 and assuming yH+ yH20Br+.The third-order rate constant k 3 ( ~ , 0 ~may ~ + ) be converted into a second-order rate using the relationship k2(HzOB~+)= k3(H20Br+)/Kf(HzOBr+)
Flgure 11. Dependence of pseudo-first-order rate constant at [Cl-] = 0 (intercept in Figure 10) upon the square of [HOCI],. Error bars represent 95% confidence interval (not shown if smaller than size of symbol).
A rate experiment was conducted in distilled HOCl solution to evaluate the dependence of kT upon [HOCl]. The data indicated in Figure 12 are consistent with a single third-order rate term of the form -d [RH]/dt = k3(CI20)[HOC1I2[RH]
(26)
Statistical evaluation of k3(cl2O)gave 0.36 f 0.01 M-'s-'. The possibility of chlorination by HOCl was evaluated by inclusion of the second-order term k2(Hocl)[HOC1] [RH] into the regression analysis. The value for k2(HOCl) obtained [k2(HOCl) = (-0.2 & 6) X M-l s-l] was small however, indicating that HOCl is an insignificant chlorinating agent under these conditions.
Discussion Formation of the electrophilic reactants was not rate limiting as indicated by the absence of terms zero order in [RH]. Zero-order terms have been observed with substrates of higher activation such as phenol and anisole and with Clz (2) and ClZO(6) as the electrophiles. Under the conditions of this study, all rate terms were first-order in [RH]. 1054
Environ. Sci. Technol., Vol. 22, No. 9, 1988
(27)
However, the k2(H OBr+) value of (1.33 f 0.01) x 1014 M-l s-l obtained from h i s data exceeds the maximum possible (diffusion limited) reaction rate (7 X lo9M-l s-l) by several orders of magnitude (3). Therefore, a bimolecular reaction with H20Br+must be ruled out. Chloride-catalyzed brominations were observed to be first-order in HOBr, H+, C1-, and RH (eq 22). Such a fourth-order term is consistent with a mechanism involving reaction between rapidly formed BrCl and RH according to Scheme I. Thus, this fourth-order term can be reduced to a second-order term by considering (from eq 6) that The rate law [HOBr][H+][Cl-] = Kh(BrCl)[BrCl]/yH+ycl-. for bromination by BrCl, therefore, is -d[RHl /dt =
(k4(BrCl)Kh(BrCl)/YH+~Cl-)[BrC11[RH]
(28)
where yH+and ycr are activity coefficients for H+ and C1at 0.1 M ionic strength. The activity coefficients YH+, ycl-, and yBr-were calculated according to the extended Debye-Huckel limiting law (28) as 0.83, 0.76, and 0.76, respectively. From eq 28 the second-order rate constant k2(BrC1) is obtained as 540 f 9 M-l s-l. The same argument is followed to reduce k4(Brz) and k4(c12)to the specific secrespectively. In ond-order rate constants kz(Br3 and k2(c12), analogy to eq 28, we may write k2(Brz) = k4(Brz)Kh(Brz) /YH+YBr(29) and obtain kz(Brz) = 0.251 f 0.004 M-' Similarly h2(Clz)
=
S-'.
k4(Clz)Kh(Cl,)/YH+YC1+
(30)
from which we obtain k2(clZ)= 58 f 9 M-' S-'. Second-order terms in HOCl may be interpreted as reactions with ClZO (6). A second-order rate constant is obtained by considering [HOCl]' = Kh(Cl2O)[c120](eq 11). Thus, b(cizo)is kZ(C120) = k3(C1~0)Kh(C1z0) (31)
Table IV. Rate Constants and Relative Reactivities for the Reaction of p-Xylene with All Possible Halogen Carriers in Aqueous HOC1-Br- Systems at 20 "C and I = 0.1 M species
rate constant"
H20Br+ BrCl ClZ ClZO Br2 HOBr HOCl
1330 10 M-2s-' 540 f 9 M-'s-l 58 f 9 M-'s-l 54 4 M-'s-l 0.251 0.004M-'s-' (1.1 X lo4 0.7) X 10"' M-'s-l
*
**
relative reactivity
b 2.2 x 2.3 X 2.2 x 1 4.4 X
103 lo2 102
lob4
NO
NO
a Error given reflects only statistical error in pseudo-first-order rate constants used to calculate second-order constants. *Relative reactivity cannot be given because bromination with H20Br+is not a bimolecular reaction.
Two sets of data are available for evaluating k2(cl20):(i) kT obtained as a function of C1- at constant [H+] extrapolated to [Cl-] = 0 and (ii) kT determined as a function of [HOC11 in distilled HOCl in the absence of C1-. The first case (eq 25) yields kzcclzo)= 54 f 4 M-' s-l, and the latter (eq 26) yields k2gzo)= 101 f 2 M-l s-l. The value obtained in distilled HOCl is higher, perhaps due to some catalysis by traces of HCl. The extrapolated value of k2(c~,0) = 54 f 4 M-l s-l is considered the more accurate one and will be used for comparison of the electrophilicity of the halogenating agents and for modeling mixed systems (9). However, the uncertainty in k2(c120) is difficult to assess because of uncertainty in Kf(C1zO) (19). Moreover, the specific reactivity determined for C120 may also be a function of the proton activity which was not evaluated in this study. Rate data and relative reactivities are summarized in Table IV. The second-order rate constants decrease in the order BrCl > C12 Cl2O> Br2> HOBr > HOC1. This order of reactivity decrease may be rationalized by invoking the electronegativity of Y (C1> Br > OH) and the polarizability of X (Br > Cl). For electrophiles which have a common Y, the brominated species is more reactive than the chlorinated analogue BrCl > ClC1, which may be explained by the polarization of the Br6+-Clh bond, and BrOH > ClOH, which may be explained with the greater polarizability of Br as compared to Cl(11). The observed higher reactivity of C1, over Br2agrees with data evaluated by the selectivity relationship (29). Previous data (5) suggest that Br2 is 1.1 X lo3 times more reactive than HOBr when p-anisate is used as the substrate, which is in good agreement with the data of this study. Activation of the aromatic system of p-xylene by the two methyl groups was sufficient to permit reaction with HOBr but insufficient to promote an appreciable reaction with HOC1. The difference in reactivity between HOBr and HOCl has been attributed by Shilov (30) to the stronger tendency of HOBr to form a polarizable encounter or a-complexwith the aromatic compound and, therefore, more facile dehydration of HOBr by H+ by yield the Wheland intermediate. The specific reactivity of H20Br+cannot be compared directly with the other halogenating species because bromination with H20Br+is not a simple bimolecular reaction. However, H20Br+has been reported to be more reactive than BrCl [Shilov and Kaniaev (31)cited by de la Mare and Ridd (I)]. The solution conditions where both species are equally significant may be evaluated by setting the rate terms for BrCl and H20Brf equal (eq 22): k4(BrCl)[c1-l = b(HzOBr+) (32) It follows that for [Cl-] = 1.23 X M, the contribution
-
of the two species to the overall bromination rate will be equal. For [Cl-] > 1.23 X lo4 M, the contribution of BrCl to the overall bromination rate will be more significant. The fastest rate observed [ k 2 p , = 540 M-l s-l] is slower than the formation of the encounter complex between C12 and phenol [2.3 X lo4 M-l s-l (S)], and therefore, formation of the encounter complex [BrCl, RH] is not likely rate limiting. This, along with the observed first-order dependence on [RH], suggests that formation of the Wheland intermediate (11) is probably the rate-limiting step. Although in many respects the chlorine and bromine systems exhibit similar chemistries, interesting differences were noted. H20C1+and Br20 appear to have no kinetic significance compared to their counterparts H20Br+and C120,respectively. No kinetic evidence was obtained for the occurrence of significant amounts of the mixed dihalomonoxide BrOC1.
Summary and Conclusions Experimental rate laws for the bromination and chlorination of p-xylene were developed in aqueous solutions of HOBr and HOCl containing C1- and Br- in the pH range 2.5-7.5, at 20 "C, and with 0.1 M ionic strength. p-Xylene was used as the aromatic substrate. Second-order rates (first order in the halogenating agent and first order in p-xylene) were derived for BrC1, C12,Cl2O,Br2,and HOBr. The specific reactivities decreased in this order (Table IV). The reactivity of HOCl was unmeasurably small. Bromine chloride has been identified as the most reactive halogenating agent; 9 times more reactive than C12 and 2.2 X lo3 times more reactive than Br2. In agreement with previous studies (321, Cl2Owas found to be a powerful electrophile. The specific reactivity of Cl2O was -2.2 X lo2 times that of Br2 and approximately as high as that of C12. This may be a significant finding because the possible involvement of C120 species in aqueous chlorination reactions is often overlooked. Reliability of this comparison depends on the accuracy of the hydrolysis constant of C120that will need to be confirmed (19). In contrast, the reactivity of HOCl was not measurable and was insignificant in the systems studied. Hypobromous acid was approximately 2300 times less reactive than Br2. For H20Brf no specific second-order rate constant can be defined because the mechanism is not elucidated in sufficient detail (3). The experimentalthird-order rate law (first order in H+, first order in HOBr, and first order in p-xylene) suggests a highly reactive H+-HOBr-RX complex (3). No kinetic evidence was obtained for the existence of significant amounts of H20C1+,Br20, and BrOC1. Registry No. H20Br+, 84906-83-2;BrC1, 13863-41-7;Clz, 7782-50-5;Cl@, 7791-21-1;Br2, 7726-95-6;HOBr, 13517-11-8; HOC1, 7790-92-3;2-bromo-1,4-dimethylbenzene, 553-94-6;p xylene, 106-42-3;2-chloro-l,4-dimethylbenzene, 95-72-7.
Literature Cited De la Mare, P. B. D.; Ridd, J. H. Aromatic Substitution; Butterworths: Scientific: London, 1959. Soper, F. G.; Smith, G. F. J. Chem. SOC.1926,1582-1591. Gilow, H. M.; Ridd, J. H. J. Chem. SOC., Perkin Trans. 2 1973,1321-1327. De la Mare, P. B. D.; Ketley, A. D.; Vernon, C. A. J. Chern. SOC.1954, 1290-1297. Derbyshire, D. H.;Waters, W. A. J. Chem. SOC.1950, 564-573. Swain, C. G.;Crist, D. R. J. Am. Chem. SOC.1972, 94, 3195-3200. Reinhard, M.; Stumm, W. In Water Chlorination: Environmental Impact and Health Effects;Jolley, R. L., et al., Eds.; Ann Arbor Science: Ann Arbor, MI, 1980; Vol. 3,pp 209-218. Environ. Sci. Technol., Vol. 22, No. 9, 1988
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(8) Grimley, E.; Gordon, G. J.Phys. Chem. 1973,77,973-978. (9) Vourdrias, E. A.; Reinhard, M. Environ. Sci. Technol.,
following paper in this issue.
(10) Reinhard, M.; Goodman, N.; Mortelmans, K. E. Enivron. Sci. Technol. 1982, 16, 351-362. (11) March, J. Advanced Organic Chemistry, 3rd ed.; Wiley Interscience: New York, 1985; pp 447-511. (12) Brown, H. C.; Stock, L. M. J. Am. Chem. SOC.1957, 79, 5175-5179. (13) Brown, H. C.; Stock, L. M. J. Am. Chem. SOC.1962,84, 3298-3306. (14) Brown, H. C.; McGary, C. W. J. Am. Chem. Soc. 1955, 77, 2310-2312. (15) Farkas, L.; Lewin, M.; Bloch, R. J . Am. Chem. SOC.1949, 71, 1988-1991. (16) Lewin, M.; Avrahami, M. J. Am. Chem. SOC.1955, 77, 4491-4499. (17) Haag, W. R. Water Res. 1981, 15, 937-940. (18) Kanyaev, N. P.; Shilov, E. A. (1938) cited in Ksenzenko, V. I.; Stasinevich, D. C. Chemistry and Technology of Bromine,Iodine and Their Compounds;Chimia: Moscow, USSR, 1979 (in Russian). (19) Reinhard, M.; Redden, G. D.; Voudrias, E. A. Water Chlorination: Environmental Impact and Health Effects; Jolley, et al., Eds.; Ann Arbor Science: Ann Arbor, MI, 1988
(in press). (20) Liebhafsky, H. A. J. Am. Chem. SOC.1934,56,1500-1505.
(21) Connick, R. E.; Chia, Y. T. J. Am. Chem. SOC.1959, 81, 1280-1284. (22) Morris, J. C. J.Phys. Chem. 1966, 70, 3798-3805. (23) Roth, W. A. 2.Phys. Chem. Abt. A 1929, 145, 289-297. (24) Shilov, E. A.; Vainshtein, F. M.; Yasnikov, A. A. Kinet. Katal. 1961, 2, 214. (25) American Public Health Association Standard Methods for the Examination of Water and Wastewater;American Public Health Association: Washington, DC, 1980. (26) Burgoyne, E. E.; Klose, T. G.; Watson, D. K. J.Org. Chem. 1955,20, 1508-1512. (27) SAS User's Guide: Statistics; SAS Institute: Cary, NC, 1982. (28) Stumm, W.; Morgan, J. J. Aquatic Chemistry, 2nd ed.; Wiley: New York, 1981. (29) Stock, L. M.; Brown, H. C. Adv. Phys. Org. Chem. 1963, 1, 35-154. (30) Shilov, E. Bull. SOC.Chim. Fr. 1963, 30, 2903-2909. (31) Shilov, E. A.; Kaniaev, N. C. R. Acad. Sci. URSS 1939,24, 890. (32) March, F. D.; Faraham, W. B.; Sam, D. J.; Smart, B. E. J. Am. Chem. SOC.1982,104,4680-4620.
Received for review May 1,1987. Accepted March 25,1988. This paper was supported by Grant CEE-81-17561from the National Science Foundation.
A Kinetic Model for the Halogenation of p-Xylene in Aqueous HOC1 Solutions Containing CI- and BrEvangelos A. Voudriast and Martin Reinhard" Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, California 94305
The general rate law for the disappearance of P-xYlene [RH] in aqueous HOC1-Br- solutions
was generally attributed to the action of hypobromous acid (HOBr) produced by the very fast and irreversible oxidation of Br- by hypochlorous acid (HOC1) according to HOCl
where k3(H*OBr+)tk2(BrCI 7 kZ(HOBr), k2(Cl& and h2(C1 0) are specific reaction rates $or H20Br+,Br 1, HOBr, Cf,, and Cl2O, respectively, was verified. Experimental disapDearance rates of RH and formation rates of RX (X = C1 br Br) were found to agree with those computed from a system of simultaneous mass balance, equilibrium, and rate equations by using available thermodynamic and kinetic constants. In aqueous HOCl solutions containing Br- and C1-, bromination of RH is catalyzed by H+ and both Brand C1- ions. Due to the superior electrophilic power of BrCl and H20Br+,the initial bromination rate is higher than the initial chlorination rate even at relatively small Br- concentrations, The rate of Br- oxidation to HOBr was not found to be rate limiting.
Many investigators have reported predominant formation of brominated over chlorinated products during chlorination of waters containing bromide ions (Br-) (1-5). The presence of halogenated (particularly brominated) organics in water is of concern, since some have been shown to be mutagenic (6). The formation of brominated organics +Presentaddress: School of Civil Engineering,Georgia Institute of Technology, Atlanta, GA 30332. 1056
Environ. Sci. Technol., Vol. 22, No. 9, 1988
(1)
where h = 2.95 X lo3 M-l s-l a t 25 "C (7). Bromides are ubiquitous water contaminants. The Brconcentrations range from 0.03 to 2.2 mg/L in groundwater, from 0.004 to 0.08 mg/L in five major U.S. rivers, and from 0.06 to 0.8 mg/L in the lower part of the River Rhine (8). In coastal zones, where seawater intrusion into groundwater may occur, and in estuarine waters, the concentration may be significantly higher due to the high concentration of Br- in seawater (65 mg/L). The frequently observed predominance of brominated over chlorinated products formed during water chlorination indicated the presence of some brominating agents, such as bromine chloride (BrC1) and bromine (Br2),that were much more reactive than HOBr (5). For example, when C1- is present in water containing HOBr (such as the C1- produced according to eq l),BrCl is formed: HOBr
Introduction
+ Br- A HOBr + C1-
+ H+ + C1- 3BrCl + H 2 0
(2) with Kf = 3.16 x lo4 M-2 a t 20 "C (9). BrCl has been shown to be a powerful brominating agent (10). Most of the published work on the formation of brominated products under water treatment conditions has been qualitative only, and the indentity of the halogenating agents have not yet been established. In a previous paper ( I O ) , we reported the specific reactivities and provided kinetic evidence for the presence of the following halogenating agents (in order of decreasing reactivity): hy-
0013-936X/88/0922-1056$01.50/0
0 1988 American Chemical Society
