Environ. Sci. Technol. 1983, 17, 261-267
(31) Duffy, J. German Patent 2 162 286; Chem. Abstr. 1972,77, 12621f. (32) Gamrath,H.; Craver, J. U.S. Patent 2 557 089; Chem. Abstr. 1951,45, 10668e. (33) Noller, C.; Dutton, G. J. Am. Chem. SOC.1933, 55, 424. (34) Foxton, A.; Jeffrey, G.; Vogel, A. J. Chem. SOC.A 1966,249. (35) Vilyanskaya, E.; Kirichenko, I.; Raxarenova, M. Zh. Obshch. Khim. 1969, 39, 2262; Chem. Abstr. 1970, 72, 43018r. (36) Kreysler, F. Chem. Ber. 1885,18, 1700. See ref 29 (lst), 1933, 6, 482.
Reference 8, pp 343-347. Nakanishi, K. “Infrared Absorption SpectroscopyPractical”; Holden-Day: San Francisco, 1962; p 56. Fieser, L.; Fieser, M. “Reagents for Organic Synthesis”; Wiley: New York, 1967; Vol. 1, p 415. Sandulesco, G.; Girard, A. Bull. SOC.Chim. Fr. 1930,47, 1300.
Fadia, M.; Shukla. V.;h’rivedi,J. J . Ind. Chem. SOC.1955, 32, 117.
Hoaglin, R.; Kubler, D.; Leech, R. J. Am. Chem. SOC.1958, 80, 3069.
Beilstein’s “Handbuch der Organischen Chemie”,4th ed.; Boit, H. G., Ed.; Springer Verlag: New York, 1966; Vol. 6, p 1261.
Garrett, K. French Patent 1 502 426; Chem. Abstr. 1968, 69, 87744s.
Received for review October 13, 1981. Revised manuscript received September 27,1982. Accepted January 3,1983. Financial support for this work was realized from Health and Welfare Canada, through Contract 683-1980181.
Ozonation of Bromide-Containing Waters: Kinetics of Formation of Hypobromous Acid and Bromate Werner R. Haag and Jurg HolgnQ” Swiss Federal Institute for Water Resources and Water Pollution Control, EAWAG, 8600 Dubendorf, Switzerland
Ozone oxidizes Br- under water treatment conditions to form HOBr. HOBr reacts further with 03,but only in its ionized form, OBr-. OBr- is oxidized not only to Br03but also to a species that regenerates Br-. The results are consistent with the following scheme of reactions: 0,iBr-
+ OBr203 + OBrO3
-+ - + - + ki
O2
OBr-
(1)
202
Br-
(2)
202
Br03-
(3)
k2
ka
where k, = 160 f 20 M-l s-l, k2 = 330 f 60 M-l s-l, and k3 = 100 f 20 M-l s-l at 20 OC. Thus, a catalytic decomposition of 0,via reactions 1 and 2 is observed. The maximum intermediate HOBr concentration is greater the lower the pH. In the presence of organic matter, HOBr reacts to form bromo organics. Thus,more bromoform was produced with humic acid at pH 6.1 than at pH 8.8. The range of conditions conducive to haloform formation is narrower than during chlorination. Many types of water that are subject to ozonation contain some bromide. For example, 0,is presently being applied to seawater for shellfish depuration in France and Spain (1)and is being considered for use as an alternative to chlorine at coastal power plants for cooling-system biofouling control in the US. (2, 3). In these waters the chemistry of O3is dominated by its reaction with Br-, due to bromide’s relatively high concentration (65 mg/L) and reactivity compared to other seawater components (4,5). Although typical Br- concentrations for drinking waters are considerably less (0-2 mg/L for groundwaters (6, 7) and 0 . 0 . 8 mg/L for surface waters (8-11)), reaction with Br- during O3 treatment must often still be considered. Previous authors have shown that O3 behaves similarly to chlorine in that it reacts with Br- to produce hypobromous acid (4,12-14). However, unlike chlorine (15,16) O3 can oxidize HOBr further to produce bromate at a significant rate even in dilute aqueous solution (4, 14). Still, the kinetics and mechanism of the latter reaction have never been studied in detail. Since HOBr can partake in bromine substitution reactions with organic solutes to produce potentially toxic brominated organics (e.g., bro0013-936X/83/0917-0261$01.50/0
mophenols and bromoform (6, 17-19)) and since HOBr may interfere with the analysis of residual Os, it is of interest to know how long and in what concentrations HOBr may exist as an intermediate in ozonated, bromide-containing waters. Moreover, it is desirable to know to what extent these reactions influence the stability of 03.
In order to answer these questions, we have determined the rate constants for the reactions of 0,with Br- and with HOBr, along with their dependence on pH. In the process, we have discovered a third significant phenomenon, namely, the reduction of active bromine (HOBr/OBr-) to Br- as a result of ozonation. Experimental Section
Reagents, Bromide-free HOBr solutions were prepared daily by vacuum distillation of aqueous Br2/AgN03mixtures at room temperature into a flask containing a few milliliters of ice-cooled 0.5 M phosphoric acid (20). Ozone water was prepared by passing 0,-containing oxygen through redistilled, ice-cooled water as described previously (21). All other reagents were analytical grade; the humic acid was a Fluka AG product. Titrisol buffers (Merck) were used to calibrate the pH electrode. All aqueous solutions were prepared with redistilled, deionized water. Analyses. Ozone was determined in aqueous solution by its UV absorbance at 258 nm ( E 2900 M-l cm-l) (21). Bromide concentrations were measured by using a bromide selective electrode (Ag/AgBr, Metrohm AG). Active bromine (HOBr,) and bromate were determined by successive iodometric titration at pH 4 and then at pH 1 (25). The measured HOBr concentrations were verified in one experiment by the UV absorbance of the dimethylamine derivative (N-bromodimethylamine, E , , , ~ ~ nm 490 M-l s-l (20)). In another experiment, the BrO,- concentrations were verified by differential pulse polarography (4) -1.56 V in 3 M NaCl at pH >7). Prior to the analysis of these bromine species, any remaining O3was either removed by a 30-s air purge or selectively destroyed by the addition of 1 mL of 8 mM dimethylamine per 20 mL of sample. Both methods required a pH above 6, to prevent purging of Br2 or slowing of the 0,-dimethylamine reaction, respectively. Bromoform was determined by GC/
0 1983 American Chemical Society
Environ. Sci. Technol., Vol. 17,No. 5, 1983 281
C
the disproportionation of HOBr to Br- and Br03- (16); blank runs using air in place of O3 demonstrated that this disproportionation proceeded at a negligible rate under our conditions. A second important observation made during preliminary experiments was that for a given amount of added hypobromite (HOBr,,,), the rate of O3 consumption increased approximately by a factor of 10 per pH unit and leveled off above the pKa of HOBr (pK, = 8.8, see Figure 2). This indicated that O3reacts predominantly with the dissociated form (OBr-) but negligibly with the protonated form (HOBr). With the help of these observations and the consideration of plausible mechanisms (as will be discussed later), the scheme of reactions shown in eq 1-4 was proposed.
O3 + Br-
- + - + - +
+ OBr2 0 3 + OBr-
Flgwe 1. Product formation during continuous ozonation of an HOBr solution: pH 8.05; IO,] = 21 pm.
0,
kl
k2
kS
O2 OBr-
(1)
Br-
(2)
Br03-
(3)
202
202
O3 + HOBr (4) Considering eq 2-4, the total rate constant for reaction of O3 with hypobromite can be expressed as kOBr- = k2 + 2k3 (5) or
1
I
4
I
I
6
I
I
k ~ ~ ~ ~= ,4 tk o2 t+ 2k3) (6) where a! is the degree of dissociation of HOBr. Note that reactions 1 and 2 form a chain reaction resulting in the catalytic destruction of ozone by the Br-/ OBr- couple: I
8
10
pH
Figure 2. Rate constant for O3 consumption by HOBr,,: (0, V) experlmental values based on eq 9, with [HOBr,], in at least 10-fold excess of [O,]; (0)same as 0, but recalculated by using steady-state analysis; (B)excess O3 added to NaBr solutions and evaluated by steady-state analysis.
ECD on pentane extracts (22). Kinetic Measurements. Unless otherwise noted, all measurements were performed at 20 f 1 "C. In general, the reactions were followed directly in 1-10-cm UV cells by recording the change in O3absorbance at 258 nm as a function of time. The reactions were usually initiated by injecting an HOBr or Br- solution into a mixture of 03, buffer, and OH. radical scavenger in the cell (23). The apparatus and procedure for kinetic measurements at constant O3concentration are described elsewhere (24). Ozone was continuously bubbled through the test solution and its aqueous concentration calculated from its UV absorbance (254 nm) in the off gases, assuming equilibrium conditions. The pH was controlled by using phosphate buffers (50 mM in final solution). The ionic strength varied from ca. 0.02 (pH 2) to ca. 0.15 M (pH 8). Dissociation Constant of HOBr. The pKa of HOBr was calibrated at 20 "C for our buffer system from the variation of UV absorbance of HOBr/OBr- as a function of pH. Results and Discussion Formulation of the Reaction Model. Figure 1 gives typical results of a preliminary experiment. The continuous ozonation of an initially bromide-free HOBr solution yields not only Br03- but also considerable amounts of Br(at pH 8). This phenomenon is not simply the result of 282
Environ. Sci. Technol., Vol. 17, No. 5, 1983
(7)
Conceptual Approach. Reaction 1 could be studied independently of reactions 2 and 3 by using low pH or a large excess of Br-. On the other hand, the reaction of O3 with OBr- (sum of reactions 2 and 3) could be separated from reaction 1 by working with a large excess of added HOBr and high pH. Batch studies showed that under these two limiting conditions the stoichiometry of reaction 1and also that of the sum of reactions 2 and 3 were close to 1 mol of 03/mol of bromine species converted. This showed not only that the reactions of Br- and OBr- could be separated but also that O3 consumption by reaction 3 is not dominant. Reaction 3 could be studied independently of reactions 1 and 2 by working under the "steady-state" condition. This was attained by adding enough O3 to allow the concentrations of Br- and OBr- to equilibrate to the point where the rates of reactions 1and 2 become equal. Insofar as the OBr- concentration is concerned, the rate of Br03formation is then no longer affected by reactions 1 and 2, because these help keep the OBr- concentration constant. Nevertheless, the reaction cycle 1-2-1... continues to consume 03,resulting in a high overall stoichiometry of about 9 mol of O3lost/mol of Br03-formed. Despite this parallel O3 loss, k3 could be determined by directly measuring the O3 consumption and correlating this to the amount of Br03- formed at the same time. The value of k2 results from the difference kOBr-- 2k3. It can additionally be evaluated from the rate of 0 3 depletion at steady state with the help of the measured values of kl and k3. Determination of kl.The reaction of O3 with Br- was studied at low pH, in order to minimize interference by
[Br-l,,
the kinetics (via reaction 1). Therefore, only the data above pH 8.5, where HOBr is largely dissociated, could be analyzed by the normal methods (23,28). The data at low pH had to be evaluated by steady-state analysis (see below). The average of the 19 measurements above pH 8.5 gave kOBr-= 520 f 50 M-ls-l at 20 "C over the ranges [O3lO= 5-25 pM and [OBr-1, = 100-370 pM. This value is based on CY values calculated from the pH by using the calibrated pK, of 8.76 for HOBr. The reaction kinetics were independent of ionic strength (0.15-0.8 M NaC1). The overall, empirical activation energy based on O3 consumption measured at pH 8.2 was 60 f 10 kJ/mol over the range 10-30 "C. Assuming negligible influence from reaction 1, kOBr- = 2.7 X 1013 exp(-60000/RT) M-' s-'. The individual values for koBr- were corrected for the parallel destruction of O3 by OH- with use of kO,OH- = 60 M-l s-l (29). This correction depended on the ratio of OBf to OH- concentration and was normally less than 10% of the total value. In this exceptional case, such a correction is allowable because hypobromite acts as an efficient scavenger for ozone-produced radicals (OH., 02-)(30), which would otherwise cause a chain O3 decomposition of unpredictable length (29). Accordingly, kOBr-was found to be independent of added scavenger concentration (0-80 mM carbonate or tert-butyl alcohol). This further indicated that the daughter radicals (presumably BrO.) do not initiate a chain decomposition of 03,contrary to the conclusions of Haruta and Takeyama (25). Kinetics at Steady State. According to eq 1-3, the concentrations of HOBr,, 03,and Br0,- may be described by the following rate equations:
M
Figure 3. Pseudo-first-order rate constant, k , [Br-I,, vs. [Br-1,: (0) pH 6, data taken on a Beckman DK-PA spectrometer: (0)pH 3-7, data taken on a Kontron Uvikon 810 spectrometer: (V)pH 3.2, from measurement of Br, formation.
reactions 2 and 3, and with a greater than 10-fold excess of Br- (isolation method). Plots of log [O,] vs. time were always linear, indicating that the reaction is first order in [O,]. A plot of the pseudo-first-order rate constant, k,[Br-],, vs. [Br-1, (Figure 3) gave a slope of 0.99, suggesting kinetics that are also first order in [Br-1. At low pH, the rates of disappearance of O3 and formation of Br2 (the latter followed at 390 nm) were equaI, verifying the 1:l (-A[O,]/A[HOBr,,]) stoichiometry. The kinetics were independent of pH (pH 3-7), ionic strength (p = 0.05-0.7 M NaC104),and scavenger concentration (0-21 mM carbonate). These results are consistent with the following rate law: -d[03]/dt = +d[HOBr,]/dt
= kl[Br-][03]
(8)
where [HOBr,] = [HOBr] + [OBr-1, over the ranges [03], = 4-100 p M and [Br-1, = 0.05-2 mM. The average of the 57 measurements in this range gave k1 = 160 f 20 M-' s-l at 20 "C. A plot of In kl vs. 1/T yielded an activation energy of 37 f 4 kJ mol at pH 2 over the range 1-30 "C. Thus, kl = 6.3 X 10 exp(-37000/RT) M-l s-l. This result agrees with that recently reported by Haruta and Takeyama (It, = 170 M-l s-l at 20 "C) (25). These values appear to be more reliable than the older ones reported by Taube (26) and Garland et al. (27) (90 and 230 M-l s-l, respectively, corrected to 20 "C based on our activation energy). Determination of kHOBr,,, and koBf-.The rate by which O3 is consumed by HOBr,, a t given pH can be formulated as
Q
-d[O3I /dt = kHOBr,tot[HOBrtotl[GI
(9)
This rate law was tested by using the isolation method by following the depletion of O3 in the presence of excess HOBr,,. Plots of log [O,] vs. time were again linear. At a given pH, the pseudo-first-order rate constant, k~o~~,,~[HoBr,,], was directly proportional to the concentration of HOBr,,, over a factor of at least 4. Figure 2 shows that kHOBr,,, increases with pH according to CY, i.e. kHOBr,tot
= ffkOBr-
d[HOBr,,]/dt a
= kl[Br-][03] - a(k2 + k3)[HOBrbt][03] (12)
-d[03]/dt = k1[Br-][O3] + a(kz+ 2k3)[HOBr,,][03] (13) d[Br03-]/dt = ak3[HOBr,,l [O3]
(14)
If the rates of reactions 1and 2 are allowed to equilibrate and the system is then observed over a period of time during which the conversion of catalyst (Br-/OBr-) to Br03- is kept small, one may make the steady-state approximation that d[HOBr,,]/dt = 0. Equation 12 then yields kl[Br-I,, = 4 k 2
+ ~3)[HOBrt,tl,s
(15)
where the subscript ss refers to the steady state. Combining this with the mass balance equation [BrTlss = [Br-ls, + [HOBrtOtls, and substituting into eq 13 give
(16)
(17)
Thus
(10)
or
-d[031 /dt = &~~-[HOBr,t][03] = ko~~-[OBr-] [os] (11) whereby kOBr- is pH independent. This kinetic law is analogous to that found for many other reactions of O3 with acids (28). However, at low pH &OBr- becomes so small that even minute concentrations of Br- interfere with
Substituting the integrated form of eq 17 into eq 14 and integrating further yield [BrO3-1 = k3a([0310- [031)/ksB
(19)
for
[HOBrtotlss
(20)
Environ. Sci. Technol., Vol. 17, No. 5, 1983
263
[BrTlss
%
I
4ooL
O 3 + W
k m
3
I
I
I
200
k 02 + OBr'
[Od
130M-'i' (Approaches k l )
la:
I 4.
I
I\
100
!T
0
20
I 60
40
T I M E ,min
I C 80
Figure 4. Dynamics of O3 consumption and BrO,- formation in an HOBr solution at pH 6.5.
jBroi1[ pM
6.1 , kU :2 8 M-'S-' pH 6.5 , k,, :6 2 and 6.7M-'S-' A pH 7.2 , kss = 28 M-'S'' o pH
I t x 10-7
I 2~ 10-7
oo
*/
I 3~10-~*
~ ( [ 0 ~ l o - ~ 0 3 1 ) / k sMs2, s
-
Flgure 5. Plot of [BO3-] vs. CY([O,]~ [03])/k, for the determination of k, (see eq 19).
Determination of k3. Equation 19 could be used to evaluate It3 by measuring Br03- formation and O3 consumption simultaneously on separate portions of a single solution of HOBr at various reaction times. Figure 4 gives the results of a typical run. Such experiments were performed at pH 6-7, where steady state is attained readily because the steady-state Br- concentration is small compared to the total bromine concentration. This ensured that [BrT], = [HOBr,], (i-e., eq 20 holds). Furthermore, the conversion to Br0; was less than 15%, so that [B~T], corresponded approximately to the added amount of HOBrht. As would be expected from eq 17, the rate of O3 disappearance was found to be pseudo first order, and k, could be determined for each run on the basis of the added [HOBr,,]. A plot of [Br03-] vs. L Y ( [ O-~[03])/k, ]~ (cf. eq 19) using all points measured in the four runs is given in Figure 5. The slope of the line gives k3 = 100 f 20 M-l s-l at 20 "C. Although the scatter among the points is large, the fit to linearity is reasonable considering the assumptions made and the necessarily small amounts of Br03- formed. More complicated mathematical analysis, 264
20
O O l 0~5
TIME,min
25
Figure 6. O3 decomposition in a Br- solution showing approach to steady state (note logarithmic scale): r.1.s. = rate-limiting step; pH 7.0; [O3I0 = 200 &I[Br-1, ; = 100 pM.
Environ. Sci. Technoi., Voi. 17, No. 5, 1983
Table I. Values of h , Determined at Steady State [HOBr,otlo, lBr-10, [ 0 3 1 0 , kss, k,, PH mM mM mM M-I s-l ~ - s 1- ~a 5.2 13.5 0.33 0.24 290 5.5 4.4 0.41 0.57 330 5.7 5.0 0.41 0.89 310 6.0 4.8 0.25 1.6 280 6.2 7.4 0.31 2.8 360b 6.7 7.2 0.26 8.2 280b 7.0 0.10 0.16 14 320 7.0 0.31 0.24 14 320 7.3 0.21 0.32 32 430 7.4 0.21 0.33 44 380 7.6 0.21 0.31 60 380 7.9 0.21 0.36 79 300 350 8.1 0.21 0.29 110 8.2 0.21 0.23 140 400 8.4 0.21 0.29 170 350 8.4 0.15 a.16 160 300b 8.5 0.14 0.21 170 310 Calculated from eq 22. Average of four runs. designed to account for the changing HOBr concentration, did not greatly improve the linearity. The slight curvature possibly indicates the formation of low concentrations of intermediates between HOBr and Br03-. Figure 5 also suggests a very slight additional pH dependence that has not been taken into account. Determination of kz.Evaluation from Steady-State Experiments. According to eq 2, the rate of Br- formation from OBr- can be formulated as follows: +d[Br-]/dt = -d[03]/dt = k2[OBr-][03]
(21)
Rearranging eq 18 gives
from which k2 could be determined on the basis of the values of kl and k3 from above and measurements of k,,. The experimental approach was to add enough O3 to HOBr or Br- solutions to achieve steady state and then evaluate k,, from the linear portions of the O3 depletion curves (cf. Figure 6) using eq 17. Normally, an excess of O3 was needed, but in the case of HOBr below pH 7, an excess of HOBr could be present. Under the former conditions steady state was often reached only after a relatively large O3 consumption (cf. Figure 6), whereas in the latter case, steady state was attained during the first 10% of O3 conversion because of the low concentration of Brrequired. A summary of the results is given in Table I. In contrast to the results obtained in experiments for determining kl, a pH dependence is now observed for the rate of O3 dis-
appearance in the presence of Br-. This indicates that in the cycle Br--OBr--Br- the reaction with OBr-, not Br-, is actually rate limiting. Furthermore the excellent agreement between the data obtained by starting with Brand with HOBr is good evidence that a catalytic process is involved, since one can enter the cycle with either species. The average value is k2 = 340 f 50 M-' s-'. Evaluation of k2 Based o n kOBr-. k2 may also be calculated by rearranging eq 5: k 2 = kOBr- - 2k3 (23) Using koBi = 520 f 50 M-l s-l (pH M.5) and k3 = 100 20 M-l s-' as determined above gives k2 = 320 f 60 M-'
*
S-1.
The good agreement between the k2 values obtained by the two methods, which are independent except for the common dependence on k3, provides further confirmation of the model reaction system. Averaging the results gives k2 = 330 f 60 M-' s-l at 20 "C over the range 1.5-370 pM OBr-. Reevaluation of kHOBr,tot. With the k2 and k3 values obtained from steady-state analysis, eq 6 gives further data for kHOBr,tot. These are plotted as solid symbols in Figure 2. They now allow interpretation of this figure in more detail: The values of kHOBr,totlevel off below pH 4. This would suggest that O3 reacts with the nondissociated HQBr, but at most very slowly. The apparent rate constant is kHOBr I0.013 M-' s-l, but this low value is uncertain due to interfering O3losses by wall effects and gas exchange. The theoretical curve in Figure 2 was calculated from kHOBr,tot = - a)kHOBr + akOBr(24) where kHoBr = 0.013 M-' 8-l and koBr- = 530 M-' S-'. It Can be seen that the use of steady-state analysis gives a better fit of the data to the theoretical curve than is obtained by assuming that only the reactions of O3with OBr- are involved (empty symbols). The difference of a factor of about 2 corresponds to the added O3 consumption by reaction 1. Mechanisms. Oxidation of Br-. Haruta and Takeyama (25) have discussed the mechanism of this reaction and concluded that the formulation given in eq 1 is most consistent with the data. Reduction of OBr- to Br-. That a reduction of active bromine as formulated by reaction 2 occurs is supported by analogy to other halogen reactions. We have observed the same phenomenon with chlorine, i.e., the formation of Cl- from HOCl/OCl- as a result of ozonation (31). We propose the following mechanism for such reactions:
O3 + OX-
slow
fast
O2 + XOO-
(25)
+
xoo0 2 x(26) This is reasonable because O3is a good oxygen atom donor and halogen peroxides (XOO-) are known species. They are formed by the reaction (eq 27) of H202with hypohalites H202 + OX- H2O + XOO(27) (32)and decompose rapidly to O2and halide as in reaction 26. In the case of the H202/bromine system, the Brproduced may be reoxidized to HOBr/OBr-, leading to a catalytic destruction of H202 (33, 34). Thus, with the exception of Br03- formation, the 03/bromine system is analogous to the H202/brominesystem both in observed phenomena and in mechanism. It should be noted that the reduction of active bromine to bromide as observed in this study cannot be due to a preliminary decomposition of O3 to such reducing species
as H202or O p because the maximum possible amount of reducing agent formed by O3 decomposition is not nearly sufficient to account for the large amounts of Br- formed. Taking the conditions in Figure 1 as an example, it may be calculated that, at the very most, 2 pM Hz02could be formed from O3 in the first 20 min, whereas at least 60 pM Br- was produced in this time. In a formal sense, one should not consider the formation of Br- a reduction at all but rather a "disproportionation", because the oxygen of the hypobromite ion is being oxidized at the same time. In other words, the reaction proceeds as an oxidation, followed by a decomposition of the peroxide formed. Oxidation of OBr- to Br03-. Of all the possible mechanisms for Br03- formation, we have used the following in the model: O3
+ OBr- -!% O2 + BrO;
O3 + Br02- -% O2+ Br03-
(28) (29)
This was chosen based partly on its simplicity and partly on its similarity to the oxidation of HOBr to Br03- by HOC1, where bromite is known to be an intermediate (16). It is also consistent with the fact that the reaction of O3 with Br02- has a rate constant (k3J greater than lo6 M-' (35),which prevents Br02- from building up to measurable concentrations. However, this mechanism is probably too simple. Richardson et al. (4)have presented a more complete mechanism by assuming the likely intermediates BrO. and Br02,but they have ignored direct reactions with O3which our results show to be important. We were unable to detect Br02 by its UV absorbance (emar 1000 M-' cm-l at 475 nm (36))in our solutions, but this does not preclude it from being a kinetically important intermediate. Indeed, pure Br02 may be synthesized by the reaction of Br2with O3 (37). Nevertheless, it is difficult to write a reasonable mechanism with an overall stoichiometry significantly different from two O3 molecules per Br03- formed, and thus the validity of the above mechanism should not affect the interpretation of the kinetic data greatly.
Practical Applications Continuous Ozonation of Br--Containing Waters. Figure 7 gives examples of the variation of bromine species concentrations as a function of time under drinking water treatment conditions (- 1 mg/L 03,pH 7-9). The following results should be noted: (1) The sum of the analytically measured species [Br-] + [HOBr,,] + [BrO
