Environ. Sci. Technol. 1994, 28, 1234-1242
Bromate Formation during Ozonation of Bromide-Containing Waters: Interaction of Ozone and Hydroxyl Radical Reactions Urs von Gunten* and Jiirg Hoigne Swiss Federal Institute for Water Resources and Water Pollution Control, EAWAG, 8600 Diibendorf, Switzerland
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Kinetic simulations have been tested by laboratory experiments to evaluate the major factors controlling bromate formation during ozonation of waters containing bromide. In the presence of an organic scavenger for OH radicals, bromate formation can be accurately predicted
by the molecular ozone mechanism using published reaction rate data, even for waters containing ammonium. In the absence of scavengers, OH radical reactions contribute significantly to bromate formation. Carbonate radicals, produced by the oxidation of bicarbonate with OH radicals, oxidize the intermediate hypobromite to bromite, which is further oxidized by ozone to bromate. During drinking water ozonation, molecular ozone controls both the initial oxidation of bromide and the final oxidation of bromite. OH radical reactions contribute to the oxidation of the intermediate oxybromine species. Bromate formation in advanced oxidation processes can be explained by a synergism of ozone and OH radicals.
Introduction Bromate is a disinfection byproduct (DBP) of the ozonation of bromide-containing waters. Its carcinogenicity in animal experiments has created an interest in the regulation of this compound in drinking water treated with ozone (1). The discussion of a regulatory standard for bromate in drinking water is governed by three main factors: (i) toxicity of bromate, (ii) detection limit in a natural water matrix, and (iii) disinfection safety criteria (2).
The bromate concentration in drinking water associated excess lifetime cancer risk of 10-5 is 3 yug/L based on a linearized multistage model for a consumption of 2 L/day by a 70-kg adult (3). Based on the present analytical feasabilities, the World Health Organization (WHO) recommended a provisional guideline value of 25 ng/L of drinking water (3). This decision may postpone the actual bromate discussion temporarily, because only waters with high bromide levels (>200 /xg/L) will be of concern for standard ozone treatments under these circumstances (4). However, the tolerable bromate concentration will likely be subject to further discussions, and some basic studies are required to identify possible procedures for the minimization of bromate production in drinking water treatment. Bromide concentrations in raw waters for drinking water production vary from a few micrograms per liter up to several milligrams per liter. In coastal groundwaters, elevated bromide concentrations are frequently attributable to the infiltration of ocean water. The major sources of bromide in inland waters are related to local geological situations, natural fractionation, and anthropogenic bromide immissions as soda production, potassium, and coal
with an
*
Author for correspondence.
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Environ. Sci. Technol., Vol. 28, No. T, 1994
mining. A compilation for varied water published in ref 4.
resources
is
Chemical Reactions
Molecular Ozone Mechanism. The reactions controlling the well-known mechanism for direct interaction of molecular ozone with bromide are given in Table 1. These data
can be used to predict bromate formation as function of bromide concentration, pH, and ammonia concentration. A schematic representation of the molecular ozone mechanism is given in Figure 1. Because reaction 1 is the only reaction for which the activation energy is known [37 ± 4 kJ/mol (8)], it is difficult to predict the behavior of the whole reaction system at temperatures other than 20 °C, the temperature at which the rate constants were determined. In pilot plant experiments performed on surface waters, a 20% increase of bromate was found for an increase of 10 °C (10). In a previous study, we performed a kinetic simulation to predict bromate formation by molecular ozone as a function of its decisive parameters (4). An example of a
a
model calculation accounting for reactions 1-5 is presented in Figure 2. Ozonation transforms bromide efficiently into hypobromous acid and hypobromite. For the rate constants listed above, 77 % of the hypobromite formed then reacts through reaction 2 back to bromide and 23 % through reaction 3a to form bromate. The oxidation of hypobromous acid by ozone is very slow (reaction 3b) and, therefore, does not contribute significantly to bromate formation. Influence of Ozone Concentration. Changes in water treatment conditions can alter the rates of ozone decomposition and other reactions significantly. To have a standardized measure for the ozone concentration (c) acting during a reaction time (t) we defined the ozone exposure as the ct value [mg/L*min]. This is represented by the integral under the ozone depletion curve in a plot of ozone concentration vs time as shown in Figure 3. Bromate formation by molecular ozone proceeds through reactions that are all first order in ozone concentration. It is thus possible to make standardized plots of bromate concentration as a function of the ozone exposure characterizing the ozonation process. Influence of pH. Whereas reaction 1 is independent of pH, the overall rates of reactions 2 and 3 decrease with decreasing pH due to the masking of hypobromite through protonation: the pKa of hypobromous acid (discussed below) is approximately 9 (eq 5). In the pH range of 7-8, only 1-10% of [HOBr]tot (in the form of OBr~) takes part in reactions with molecular ozone. Influence of Ammonia. If ammonia is present, hypobromite is masked by the formation of monobromamine (eq 6). In the pH range of drinking water treatment, this reaction is relatively fast. Monobromamine is slowly oxidized by ozone to nitrate and bromide. After most of the ammonia is depleted, bromate is formed according to reactions 1-5. Thus, ammonia causes a lag time for the 0013-936X/94/0928-1234$04.50/0
©
1994 American Chemical Society
18
Reactions and Rate Constants for Molecular Ozone Mechanism
Table
1.
k or p.Ka (20 °C)
reaction
no.
Br 02 + OBr OBr20a + BrOBr- —*• BrOa" + 02 BrOa" + 02 + H+ 03 + HOBr O3 + O3 + Oa +
1
2
3a
3b
-
-»
—
BrOa" + 03—*’ Br03~ HOBr H+ + OBr HOBr + NH3 - NHaBr + H20 O3 + NH2Br -* Y° Y + 2O3 -* 2H+ + NO3- + Br- + 30a NH4+ s h+ + nh3
4 5
6 7
8 9 “
Y
are
refs
160 M-1 s-1 330 M-1 s-1 100 M_1 s-1 10s M_1
s_1
9(8.8) 8 X 107 M-1 s-1 40 M-1 s_1 kg » k9.3
5-8 5-8 5-8 8 8 8 9 9
9
unknown products that react in later reactions. Figure 3. Typical ozone depletion curve In a batch-type reactor (sample from Lake Zurich spiked with 1 mg/L of bromide; ozone dose 2 mg/L, pH 8). The area under the curve represents the ozone exposure (ct value) which is represented by the dashed line.
Br03"
Table 2. Reactions and Rate Constants for Hydroxyl Radical Mechanism reaction
no.
Brominated organic comp. Figure 1. Reactions occurring upon ozonation of bromide-containing solutions. Only reactions In the absence of OH radicals are shown (adapted from ref 8).
BrOH-
11
Br + OH
12 13 14 15 16 17 18 19 20 21 22 23
BrOH- —Br + OHBr + Br -» Bra-
24 25 26
BrOa- + OH - BrOa + OHBr02" + CO3- —*- BrOa + CO322BrOa 0.15 M. In our experiments, the ionic strength was in the range of 1-4 mM. A comparison between experimental data and model calculations at pH 7 showed that our experiments can be best modeled by setting the pKHOBr to 9 (data not shown). Ammonium. Figure 6 shows a continuous ozonation of solutions (pH 7) containing 0.42 mg/L of bromide in the absence and presence of 0.05 mg/L of ammonium. The experimental results correspond with the assumption that ammonium masks the hypobromite. The monobromamine formed as an intermediate (eq 6) is slowly oxidized by ozone to form nitrate and bromide (eqs 7 and 8). Thus, the lag time observed for the formation of bromate in the presence of ammonium is as expected from model calculations based on reactions 1-9, Even though our experiments covered a range of molar Br~/NH4+ ratios of 1-4 at pH 7, a 100% molar recovery of NH4+-N as NO3--N was Environ. Set. Technol,, Vol. 28, No. 7, 1994
1237
Ozone exposure [mg/L
x
min]
Figure 7. Comparison of measured (symbols) and calculated bromate concentrations. Continuous ozonation in the presence of 0.05 mg/L of ammonium and an OH radical scavenger (2 mg/L of octanol) at pH 7.1. Upper curve, 1 mg/L of bromide; lower curve, 0.22 mg/L of bromide.
Ozone dose [mg/L] Figure 9. Bromate formation as a function of the ozone dose. Data for complete decomposition of ozone are shown for Br (squares) and OBr (circles) as starting compounds both for pH 11 (a) and pH 10 (b). The dashed lines represent the maximum OBr concentration formed if only eq 1 and the half-life of ozone are considered. The good agreement shows that the oxidation of Br to OBr by ozone is a key reaction and that OBr is then efficiently oxidized by OH radicals. Initial bromide concentration, 1 mg/L; hypobromite as the starting compound has the same molar concentration.
Figure 8. Evolution of nitrate as a function of ozone exposure. Squares, experimental data; dashed curve, kinetic simulation based on eqs 1-9. Continuous ozonation in the presence of 0.22 mg/L of bromide, 0.05 mg/L of ammonium, and 2 mg/L of octanol at pH 7.1. The initial ammonium concentration is also shown.
achieved. This indicates that the formation of NHBr2 and NBrs leading to a conversion of NH4+-N into N2-N is not significant. Only at pH >7 the formation of NHBr2 might become an important process; this will be discussed below. In Figure 7, the evolution of bromate in a continuous ozonation experiment is shown for two bromide concentrations in the presence of ammonium. The good agreement of the measured and simulated data shows that the processes in the bromide-ammonium-ozone system in which OH radicals are scavenged (by octanol) can be accurately described by the mechanism which accounts for molecular ozone reactions (eqs 1-9). Figure 8 shows the corresponding production of nitrate resulting from the oxidation of monobromamine (eqs 7 and 8). Complete oxidation of ammonium to nitrate is observed. The evolution of nitrate calculated by the kinetic model (eqs 1-9) shows an excellent correspondence with measured data. The ozone exposure for which the formation of nitrate levels off corresponds to the end point of monobromamine oxidation. Beyond this region the ozonation process can be described by eqs 1-5. Oxidation by Hydroxyl Radicals. To investigate the contribution of OH radicals to bromate formation in solutions that do not contain organic OH radical scav1238
Environ. Sci. Technol., Vol. 28, No. 7, 1994
engers, batch-type experiments with varying ozone dosages performed at pH 10 and 11. The decay of ozone in such solutions is initiated by OH- and accelerated by a
were
chain reaction, in which OH radicals act as chain carriers (22). The OH radical yield from decomposed ozone in such solutions can be considered as nearly independent on pH and corresponds to approximately 50% of the decayed ozone (23). In these experiments, bromate was measured after the ozone decomposition was completed. At pH 11, the measured half-life of ozone was approximately 2 s. Bromate produced from bromide as a function of the ozone dose is represented in Figure 9a. Approximating the decay of ozone by first-order kinetics, it is possible to estimate the maximum amount of hypobromite formed by the oxidation of bromide with molecular ozone (reaction 1). If we assume that all hypobromite formed by this reaction is immediatly oxidized to bromate (reactions 4 and 11-27), bromate concentrations would fall on the dashed line in Figure 9a. The correspondence of the experimental results with this estimate indicates that the oxidation of bromide to hypobromite by ozone is a key reaction even when ozone is rather quickly transformed into OH radicals. The circles in Figure 9a stand for a similar experiment, where the same molar concentration of bromine was present as hypobromite. The enhanced formation of bromate in this case shows once again the importance of hypobromite as a decisive intermediate in bromate formation through the OH radical mechanism: as soon as hypobromite is formed, it can be rapidly transformed into bromate by OH radicals. Lowering the pH from 11 to 10 increases the half-life of ozone from 2 to approximately 20 s. Bromate concen-
200
-1-i-
ryzn
nb scavengerp-
^ 150
...........
o>
................................. .................. i /
=L
®
1
c..* •
0.1 mM t-BuOH 4 mM
P:--0
100
[CO|-]tot
......................................./.Cr,.................
(0
a'/ /
E
2 50
...............n......................................................0 1
QQ
k
h 1
mM t-BuOH
—
____-A.....A _I__
i
0
-O
2
3
4
5
Ozone dose [mg/L] Figure 10. Effect of OH radicals on bromate formation at pH 8. Circles: measured bromate formation in the presence of an organic OH radical scavenger (0.1 mM t-BuOH). Curve “ozone": calculated corresponding bromate formation only considering the direct reactions of ozone. Squares: measured bromate formation including OH radical reactions. Curve “OH + ozone": calculated corresponding bromate formation through simultaneous ozone and OH radical pathways. Initial conditions: 1 mg/L of bromide, 2.5 mM carbonate at pH 8, and an ozone dose of 2 mg/L.
trations measured for these conditions are shown in Figure 9b. At pH 10, there is only a small difference in bromate formed when either bromide or hypobromite are used as the starting compounds. This is due to the longer halflife of ozone, which alllows a preliminary production of hypobromite from bromide. The overall bromate formation observed in this experiment (Figure 9b) corresponds quite well with the calculated maximum concentration of hypobromite resulting from ozonation of the bromidecontaining solution (reaction 1). This is another indication that hypobromite is a key intermediate, which is then rapidly transformed into bromate by OH radicals according to reactions 11-27. These experiments at pH 10 and 11 clearly show the importance of both ozone and hydroxyl radicals for bromate formation (see Figure 4). If the pH is lowered to a typical range for drinking water treatment (pH 7-8), the relative importance of the mechanisms is subject to change because ozone decomposition is much slower. Role of OH Radical Mechanism at pH 8. To test the contribution by the OH radical mechanism, batch-type experiments with a single ozone dose (2 mg/L) were performed at pH 8 (alkalinity 2.5 mM). Bromate formation as a function of the ozone exposure in the absence and in the presence of t-BuOH (0.1 mM) is shown in Figure 10. The difference in bromate formed is due to the inhibition of OH radical reactions. The curves included in Figure 10 show bromate formation calculated for oxidation by molecular ozone and the simultaneous reaction of ozone and OH radicals. The calculations are based on the above kinetic data, assuming an OH radical yield of 50% with respect to decayed ozone (23). A firstorder rate constant k = 1.5 X 10-4 s-1 for the conversion of O3 into OH radicals was calculated from a fit of observed ozone decay curves. There is a good agreement between experimental and calculated data. For the experimental conditions considered here and for low ozone exposures, the OH radical mechanism can contribute significantly (>60%) to the bromate formation for low ozone exposures. For higher ozone exposures, the contribution of OH radicals is smaller than that of molecular ozone. The relative importance of the oxidation
Figure 11. Effect of carbonate radicals as secondary oxidants on bromate formation. Squares, no scavenger for OH radicals; triangles, 0.1 mM t-butanol (>90% scavenging of OH radicals); circles, 0.1 mM t-butanol and 4 mM carbonate (>90% scavenging of OH radicals by carbonate). Initial conditions: 0,5 mg/L of bromide, 2.5 mM [HP042“]ta, at pH 10.2.
by ozone or OH radicals is also dependent on the bromide concentration. For high bromide concentrations, the molecular ozone mechanism is more important because the steady-state concentration of OH radicals is relatively low and limits the reaction with hypobromite and hypobromous acid. For low bromide concentrations, the steadystate OH radical concentration is higher, and thus there will always be sufficient OH radicals for an effective oxidation of hypobromite and hypobromous acid. Alkalinity. To investigate the influence of carbonate on bromate formation, a series of experiments were performed at pH 10.2. A single ozone dosage in the range of 1-5 mg/L was applied to 0.5 mg/L of bromide in 20-mL batch flasks. Bromate was measured after the complete decomposition of ozone. The curve on top of Figure 11 represents the results of solutions containing a 2.5 mM phosphate buffer. In this system, OH radical scavenging by the buffer is negligible [&hpo4,oh = 1-5 X 105 M'1 s_1, (15)], and both the reactions with molecular ozone and OH radicals lead to bromate formation. When t-BuOH (0.1 mM) is added to the same solution, approximately 90% of the OH radicals are scavenged, and bromate formation decreases significantly (see Figure 11). In a third experiment, a solution containing 4 mM [C032']tot in addition to phosphate buffer and t-BuOH was used. Under these conditions, more than 90 % of the OH radicals were scavenged by [C032_]tot and transformed into COsradicals (reactions 27 and 28). In contrast to OH, CO3was not scavenged by t-BuOH but was still capable of oxidizing hypobromite (reaction 26). The result of the third experiment at pH 10.2 is consistent with this model: The measured bromate formation in the presence of carbonate is comparable to experiments performed without an OH radical scavenger (i.e., results were similar to those for the solutions containing only phosphate). This hypothesis was further tested through ozonation experiments on bromide-containing solutions at pH 8 with carbonate concentrations ranging from 1 to 10 mM (results not shown). The calculated fraction of OH radicals scavenged by [C032"]tot varied between 50% (1 mM [C032-]tot) and 90% (10 mM [C032-]tot)- No systematic variation was observed in the bromate formed for differing alkalinities. These results again support the model assumptions in eqs 16-18, where [HOBr] tot can react either Environ. Sci. Technol., Vol. 28, No. 7, 1994
1239
Table
3.
Bromate Formation in Pilot Plant (pH
=
°C) (25)'. Experimental and Calculated Data
initial Br-
O3 dose
(mg/L)
(mg/L)
0.22 1.79 2.03 3.33
8,
T
=
20
O3
2 2
3.5 3.5
Br03- (Mg/L) measd/ Br03-6 (mg/L-min) measd" calcdti calcd (Mg/L) exposure" 2.28 2.98 5.85 5.45
5
7
56
55 110 149
141
120
0.7 1.0 1.3 0.8
2
35 78 118
“ Estimated from experimental data given by Krasner et al. (25). Calculated bromate by ozone mechanism (reactions 1-9).c Measured by Krasner et al. (25). d Calculated considering reactions 1-27. 6
Figure 12. Ozonation of natural water samples (0.45-jum filtered) in batch reactors. Squares, Lake Zurich; triangles, Lake Greifensee; closed symbols, undiluted water; open symbols, 1:1 dilution of the corresponding lake water (conservation of alkalinity). The curves are equivalent to those shown in Figure 10 for an identical water without DOM. All bromate measurements fall around the calculated curve for a simultaneous ozone and OH radical pathway. All lake waters were spiked with 1 mg/L bromide, and the ozone dose was 2 mg/L.
with OH or carbonate radicals. Correspondingly, kinetic simulations based on reactions 1-27 have shown that the bromate formed through the OH radical mechanism is nearly independent of total carbonate concentrations. This is the first case reported in which carbonate radicals are observed as transient species that act as secondary oxidants in ozonated water. Organic Matter. The transformation of ozone into OH radicals is influenced by the type and concentration of dissolved organic matter (DOM) present in natural waters. A set of experiments were performed with water from Lake Zurich and Lake Greifensee spiked with 1 mg/L of bromide to study the impact of DOM on bromate formation. The experimental and the calculated data based on reactions 1-27 are compared in Figure 12. Since no reactions with DOM are included in our model, the agreement of experimental and calculated data suggests that bromate formation per ct is independent of DOM. Because the contribution of the molecular ozone mechanism is welldefined for a given ozone exposure, this means that the effect of OH radicals has to be similar for the different waters. This is notable because the second half-lives of ozone vary between 2 (Lake Greifensee) and 10 min (2.5 mM bicarbonate, no DOM). It could be expected that a factor of 5 difference in the second half-lives of ozone would lead to significant differences in the OH radical concentrations. However, it has been shown that the steadystate ratio of [0H]/[03] during continuous ozonation of several DOM-containing surface waters has a fairly constant value, averaging 1.3 X 10~7 (24). This implies that the rate of OH radical consumption by DOM is proportional to the ozone depletion rate and is also controlled by the DOM. The small influence of type and concentration of DOM on bromate formation is also a consequence of the relatively slow reaction rate between HOBr and DOM. The steadystate concentration of HOBr thus remains almost unaffected by DOM. Knowing the characteristics of a DOMfree system, it is possible to calculate total bromate formed (within 20-30%) as a function of ozone exposure for any natural water. 1240
Environ. Sol. Technol., Vol. 28, No. 7, 1994
Based on this approach, we attempted to estimate bromate formation in actual pilot plant experiments, published by Krasner et al. [(25), Table 3, therein). The results are summarized in Table 3 together with our calculations based on eqs 1-27. The kinetic simulations were performed with an identical data set as used in our experiments at pH 8. Depending on the ozone exposure (ct value) and bromide concentration, the calculated amount of bromate formed through reactions with molecular ozone contributes between 30 and 80% of the overall bromate formed (last column, Table 3). As already outlined, the OH radical mechanism is more important for short ozone exposures and small bromide concentrations, whereas for longer ozonation times and high bromide levels, the oxidation by molecular ozone contributes more to bromate formation. The agreement of measured and calculated data is within the expected uncertainty of ±2030 %. However, the bromate measurements in those pilot experiments on waters containing 2.03 and 3.33 mg/L of bromide are somewhat surprising: For the higher initial bromide concentrations and similar ozone exposures, lower bromate concentrations were found. The calculated values follow an opposite trend. A reason for this contradiction might be the uncertainty in the estimated ozone exposure, which could only be based on data given for residual ozone concentrations in ref (25). Ammonium. To test the influence of OH radicals on bromate formation in the presence of ammonium, batchtype experiments with solutions containing 1 mg/L of bromide and 0.12 mg/L of ammonium were performed at pH 8 applying an ozone dose of 2 mg/L. For such solutions without organic OH radical scavengers, measured bromate formation as a function of ozone exposure was much slower than expected from model calculations. Under these experimental conditions, the molar ratio of Br~/NH4+ of 2 might lead to NHBr2, which is oxidized more slowly than NH2Br (see above). Role of Hydrogen Peroxide. In advanced oxidation processes, often a combination of ozone and hydrogen peroxide are applied. Deprotonated hydrogen peroxide (H02-) acts as initiator for the chain reactions that transform ozone into hydroxyl radicals [(22), /?o3,ho2- = 5.5 X 106 M_1 s-1]. In our experiments (pH 8, bromide 1 mg/L, 2.5 mM [C032_]tot. ozone dose 2 mg/L), the half-life of ozone is in the range of200-2.4 s for a H2O2 concentration between 0.09 and 7.1 mg/L (2.6_21041M). Figure 13 shows the experimentally determined bromate formation as a function of the H2O2 concentration after complete consumption of ozone. With increasing H2O2 concentration, the bromate formed increases to a maximum at a H2O2 concentration of 0.7 mg/L (20 ftM) and then decreases readily with increasing H2O2.
500
r-J
400
"3>
£
300
m
200
Io
100 0
2
4
6
8
H202 [mg/L] Figure 13. Ozonation in batch reactor showing bromate formation as function of the hydrogen peroxide concentration. Solution: 1 mg/L of bromide, 2.5 mM carbonate, pH 8, ozone dose 2 mg/L (42 gM). a
As the hydrogen peroxide concentration increases, the exposure of the solution decreases, and the OH radical production is accelerated. The optimum for the synergistic effect of ozone and OH radicals with respect to bromate formation was observed at a H2O2 concentration ozone
of 0.7 mg/L. For higher concentrations of H2O2, transformations of O3 into OH radicals become too fast, and the direct reaction of O3 with Br_ or OBr- becomes less important. This leads to a decrease in bromate formation. An additional pathway leading to decreasing amounts of bromate with increasing H2O2 might be the direct reaction of hydrogen peroxide with hypobromous acid (26):
H202 + HOBr
—
H+ + Br~ + H20 + 02 k2S
=
2 X 104
M'1 s_1 (28)
The half-life for HOBr controlled by reaction 28 can be calculated to be in the range of 13-0.16 s for the concentrations of hydrogen peroxide (2.6~210 mM) applied in our experiments. However, model calculations indicate that the impact of H2O2 on the half-life of ozone, and thus the production of secondary oxidants, is far more important than its direct reaction with hypobromous acid. For advanced oxidation processes in semi-batch systems, where both ozone and H2O2 are continuosly added to a bromide-containing solution, a dramatic increase of the bromate formation was found in preliminary experiments. A more detailed discussion of the influence of advanced oxidation processes on bromate formation is given in ref 27.
Conclusions
Bromate formation by the ozonation of bromidecontaining waters can be described by a limited set of chemical equations including direct reactions with molecular ozone and reactions with OH radicals. In systems where OH radicals are scavenged by organic compounds, bromate formation can be accurately predicted as a function of pH and ammonia concentrations using kinetic calculations based on published reaction rate constants. An excellent agreement between experimental data and model predictions was achieved by slightly adjusting the reported pKa value of HOBr.
In systems where OH radicals are not scavenged by organic compounds, considerably higher amounts of bromate can be formed. The resulting bromate concentration in systems where a substantial part of the ozone is transformed into OH radicals is often 30-70% greater than the bromate concentration found in systems where only reactions with molecular ozone occur. In drinking water ozonation, an ozone exposure of 4 mg/L-min at a pH of 8.5 and an initial bromide concentration of 200 Mg/L will result in a final bromate concentration of approximately 25 Mg/L, the present standard used by the WHO. By lowering either one of the three decisive input parameters (ozone exposure, pH, or bromide concentration), the bromate concentration can be reduced below the current standard. The effect on bromate formation of the following parameters was studied: Alkalinity. Carbonate radicals, which are formed through OH radical scavenging, react as secondary oxidants for hypobromite. Ammonium. In the presence of organic OH radical scavengers, the presence of ammonium results in a lag time for bromate formation. At pH 8, the bromate standard of 25 Mg/L is reached for a ct value of 4 mg/L-min and an initial bromide concentration of 1 mg/L. The effect of ammonium on bromate formation in presence of OH radicals is not entirely understood. Dissolved Organic Carbon. DOC accelerates the transformation of ozone into OH radicals but also scavenges them. In our experiments, the OH radical yield per ozone exposure (ct), appears to be independent on DOM. Hydrogen Peroxide. Hydrogen peroxide highly accelerates the production of OH radicals through ozone decomposition. The maximum bromate production occurred in systems where ozone decomposition was slow enough to allow a synergistic effect of ozone and OH radicals. Acknowledgments
The authors thank H. Bader, M. Boiler, S. Canonica, and D. Sedlak for fruitful discussions and corrections in the manuscript.
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Received for review August 3, 1993. Revised manuscript ceived January 31, 1994. Accepted March 15, 1994. * ®
re-
Abstract published in Advance ACS Abstracts, April 15,1994.
