Ind. Eng. Chem. Res. 1995,34, 2314-2323
2314
Advanced Oxidation Processes. A Kinetic Model for the Oxidation of 1,2-Dibromo-3-chloropropane in Water by the Combination of Hydrogen Peroxide and U V Radiation William H. Glaze* Department of Environmental Sciences and Engineering, School of Public Health, The University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-7400
Yiishyan Lay Taiwan Environmental Protection Center, Central Taiwan, 480 Jan-Shin Road, Taichung, Taiwan, ROC
Joon-Wun Kang Department of Environmental and Industrial Hygiene, College of Health Science, Yonsei University, Wonju Campus, Kangwon-Do, Korea
A kinetic model for the oxidation of organics in water by the combination of hydrogen peroxide and W radiation is described. The model is based on literature values for a series of reactions initiated by the photolysis of hydrogen peroxide by W radiation into hydroxyl radicals, to which is added a term for the direct photolysis of the organic. The model is tested with data on the oxidation of a compound, 1,2-dibromo-3-chloropropane(DBCP), at low levels (12). Samples (3.0 mL) were taken into 5 mL vials containing 0.45 mL of horseradish peroxidase solution (8 unitdml), the p-hydroxylphenylacetic acid (in excess), and 0.45 mL of 0.35 M potassium hydrogen phthalate buffer (pH 5.5). After a few minutes, 0.45 mL of 0.4 M NaOH was added. Fluorescence of the solution was measured with a Perkin-Elmer model LS-3B spectrometer as described by Lazrus et al. (1985). Solutions with peroxide levels beyond the dynamic range of the method (-0.6 ppm) were diluted with laboratory grade water before analysis. The method for analyzing DBCP, modified from USEPA method 501.2, consisted of microscale liquidliquid extraction using hexane (Fisher Scientific, Optima grade) with a water-to-extractant volume ratio of 2:l. The hexane contained the internal standard DBP at a concentration of 200 pgA,. A Varian model 3500 gas chromatograph equipped with an autosampler, a 30 m x 0.32 mm 0.d. fused silica capillary column (DB-5, J&W Scientific), and a Ni63 electron capture detector were used for analysis of the extracts. The carrier and makeup gas were 99.999% nitrogen. Other GC conditions were as follows: splitless injection; injector temperature, 100 "C; oven temperature program, 50 "C (1 min), 50-110 "C at 30 " C h i n , 110 "C (2 min), 110230 "C a t 30 "Chin, 230 "C (1 min); and detector temperature, 300 "C. Quality control features as specified in USEPA method 501.2 were used in preparing method calibration standards. Procedure for Kinetic Runs. For a typical oxidation run,70 L of Arrowhead distilled water was pumped into the reactor with a pump constructed so that only poly(tetrafluoroethy1ene) (PTFE) and stainless steel contacted the water. The initial sodium bicarbonate concentration was adjusted to the desired value, and DBCP was spiked into the reactor to give a target concentration of approximately 300 ,@ In someI cases, ,. the pH of the water was adjusted using either sodium hydroxide or hydrochloric acid. After several minutes of mixing at a stirrer speed of 500 rpm, a time zero sample was collected and then an appropriate volume of the H202 stock solution was added to the reactor. Just before the W lamps were started, "time zero" samples were taken for the measurement of initial DBCP and peroxide. During the course of each run, samples for DBCP and peroxide analysis were taken into clean 40 mL vials containing three drops of a 1% sodium thiosulfate solution (to quench any active halogen that
2316 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 0
-1
h
0
F
%P x &
-2
0
m
8
E
-3
. . . . . . . . . . . . . . .
-4
0
800
400
1200
d
1600
TIME, s Figure 1. Pseudo-first-order rate plots of kinetic data for the photolytic oxidation of DBCP in 0.10 mM total carbonate solutions with 254 nm radiation (0.48 W L ) .Initial concentrations [HzOzL and - _. _ correlation coefficients for linear fit: A , 0.79 mM, 0.9987; A, 1.50 mM, 0.9999; 0 , 3.00 mM, 0.9986. ~
might have been formed). Sample vials were filled headspace-free and sealed with PTFE-lined screw caps. Extraction of DBCP samples was initiated immediately after completion of an oxidation run. Actinometric Measurements. Potassium ferrioxalate was prepared by a procedure described by Hatchard and Parker (1956). A solution of the recrystallized product (0.0068 M in 0.1 M sulfuric acid) was added to the reactor, and the W lamps were started. At suitable time intervals, 2.0 mL samples were taken from the reactor and added to a 25 mL volumetric flask containing 8 mL of 0.1 N sulfuric acid, followed by 2 mL of 0.1% 1,lO-phenanthroline and 5 mL of 0.60 M sodium acetate. The solution was then diluted to the mark and the absorbance read a t 510 nm. The concentration of Fez+ was determined by comparison with standard solutions. The values of the W intensity at 254 nm (I,)for one, two, three, and four lamps in the 70 L reactor obtained from ferrioxalate actinometry using zero-order conditions were 0.26, 0.52, 0.77, and 1.04 x einstein L-l s-l, respectively. Hydrogen peroxide actinometry was used to verify these values and to determine the effective path length of the reactor (Guittonneau, 1989). The measured values of Iofor four lamps obtained a t three different initial peroxide concentrations (4.4, 5.3, and 7.9 mM) were 0.92, 0.99, and 0.91 ( x einstein L-l s-l), respectively. At an initial concentration of [H202Io = 0.054 mM and four lamps (0.48 WL), the value of d In [HzOzYdt, and thus s-l. Using I, = 2.303IOcb,was found to be 6.76 x 1.04 peinsteins L-l s-l, a value of 2.303~b= 650 M-' was calculated. The value of the extinction coefficient
of hydrogen peroxide on a loglo basis is 17.9 M-l cm-l. Hence, the effective path length of the reactor is 15.8 cm. This is a reasonable number, given the fact that the diameter of the reactor is 15 cm. E
Results and Discussion Mechanism of the Photolysis of Hydrogen Peroxide. Several fundamental studies (Hunt and Taube, 1952; Baxendale and Wilson, 1957; Volman and Chen, 1959) have shown that the primary process in the photolysis of hydrogen peroxide produces two OH radicals, initiating a chain reaction as shown in Scheme 1. Under conditions where hydrogen peroxide is the principal absorber, its rate of photolysis is given by the eq 1:
where I, is the (monochromatic)W intensity, At is the total absorbance of the solution, and @ H ~ o is ~ the primary quantum yield for photolysis of hydrogen peroxide. Baxendale and Wilson (1957) measured the quantum yield for photolysis of hydrogen peroxide a t 254 nm as 1.00 f 0.02. Later, Volman and Chen (1959) confirmed the presumption of Baxendale and Wilson that the quantum yield of the primary process (first equation in Scheme 1)is 0.49. The value obtained by Baxendale and Wilson is due to the fact that the overall process consisting of reactions 1 , 2 or 3, and 11 in Scheme 1 decomposes two hydrogen peroxide mol-
Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 2317
TIME, 8 Figure 2. Pseudo-first-order rate plots of kinetic data for the photolytic oxidation of DBCP in 4.0 mM total carbonate solutions with 254 nm radiation (0.48 W L ) .Initial concentrations [HzOzloand correlation coefficients for linear fit: 0,0.35 mM, 0.9825; 0,0.56 mM, 0.9935; A, 1.00mM, 0.9935.
ecules per photon absorbed. Thus, although the primary process has only a quantum yield of 0.49,the quantum yield of the overall process has a value of unity. The H202/1TV process is thought to consist primarily of OH radical oxidation of organic pollutants such as DBCP (Glaze et al., 1987;deLatt et al., 1988). In water with high carbonate alkalinity typical of many groundwaters, organic substrates must compete with natural organic matter, hydrogen peroxide, bicarbonate, and carbonate ions for reaction with the hydroxyl radical. The product of the reaction of carbonate or bicarbonate with 'OH is the carbonate radical 'COS- . In hydrogen peroxide-rich solutions, the fate of the carbonate radical is presumed to be the equation below (number 12 in Scheme 1). COS-
+ H,O, -.HCO,- + HO,
If the fate of the hydroperoxy radical HOz (or its conjugate base) is disproportionation to yield hydrogen peroxide, then the net reaction of the photolysis is the same when bicarbonate is present or absent. Thus, the quantum yield for the decomposition of hydrogen peroxide should be unity in bicarbonate-containing water as well as in the pure system reported by Baxendale and Wilson. This is in accordance with the conclusions reached by Peyton and Glaze (1988)who reported work in the OdUV system, in which hydrogen peroxide is formed in situ. In addition to OH radical oxidation, organic pollutants may also be directly photolyzed in the H 2 0 W process and, in some cases, they may be oxidized by carbonate,
superoxide, or other radicals. It is instructive, therefore, to compare experimental data with models based on radical reactions. Oxidation of DBCP with HzOzNV in Distilled Water. Thirty-eight batch runs were made in which DBCP was oxidized with combinations of hydrogen peroxide and 254 nm radiation in the presence of various levels of total inorganic carbon CT and a t different pH values. Table 1 is a compilation of the data for these runs. For the convenience of the reader, some of the runs are repeated in the table so as to illustrate the effect of [HzOZI,,pH, etc. Figure 1shows the kinetic data obtained in three of these runs at pH 8.0-8.2 with four U V lamps (0.48WL) and 0.10 mM CT (run codes 3204, 3304, and 3404 in Table 1). All of the other runs are similarly linear, showing that the oxidation of DBCP in the H z O W process is psuedo-first-order; i.e., d In ([DBCPlt/[DBCP1,)ldt = k,, where [DBCPlt and [DBCP], are values of the concentrations of DBCP at times t and zero and k, is the observed pseudo-first-order rate constant for overall oxidation of DBCP. The values of k, are tabulated in Table 1. The rates of oxidation of DBCP are about 1 order of magnitude slower than for tri- and tetrachloroethylene under similar conditions (Glaze and Kang, 1989a,b),which is expected for a halogenated alkane vs halogenated alkenes (Buxton et al., 1988).In the kinetic runs shown in Figure 2 (run codes 2604, 2004, and 1904 in Table l), the level of CT is 4 d, corresponding t o 200 mg/L of alkalinity as calcium carbonate. As shown by the ordinate scales in Figure 1 vs Figure 2 and the values of KO in Table 1,the effect of carbonatehicarbonate is t o slow the reaction rate ap-
2318 Ind. Eng. Chem. Res., Vol. 34,No.
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2
3
4
5
6
7
Hydrogen Peroxide, mM
Figure 3. Effect of initial hydrogen peroxide concentration on the pseudo-first-order rate constants for the oxidation of DBCP with HzO2 and 254 nm U V radiation, CT: 0, 4.0 mM; 0 , 4.0 mM; A, 0.1 mM. UV: 0, 0.48 W/L; 0 , 0.24 W/L; A, 0.48 W L
preciably, as discussed in more detail below. A value of 4 mM CTis not an atypical alkalinity for groundwater, however. Effect of Initial Hydrogen Peroxide Concentration and CT. Figure 3 shows the effect of initial hydrogen peroxide concentration on the value of k, for two W dose rates of 0.24 and 0.48 W/L and with two CT levels a t the higher W dose. It is apparent from this data that the rate of decomposition of DBCP reaches a maximum value that is dependent upon the bicarbonate scavenger level. This is presumably due to the interplay between *OH oxidation of DBCP, which is accelerated by enhanced photolysis of H202 at higher concentrations of H202, and 'OH scavenging by carbonate, bicarbonate, and HzO2. The rate data in Figure 3 a t low carbonate alkalinity show a more pronounced maximum in the values of 12, as the initial peroxide level is increased. Under low alkalinity conditions, scavenging of OH radicals will be low, especially at low peroxide levels. Hence, the rates of oxidation of DBCP are high a t CT= 0.1mM, with k, values reaching 2.7 x s-l. This corresponds to a time for 3 log, (95%) removal of DBCP of 19 min. On the other hand, a t 4 mM CT,the corresponding treatment time would have to be approximately 50 min at a
hydrogen peroxide level twice as high. Hence, it appears that cost-effective treatment of halogenated alkanes such as DBCP would be achieved by presoftening of the water to low carbonate alkalinity levels. Effect of W Intensity. Figure 4 shows that pseudofirst-order rate constants for DBCP oxidation are proportional to the W intensity, at least under the conditions used in this work. Direct Photolysis of DBCP. The data in Table 2 show that in the absence of peroxide there is a small residual rate of decay of DBCP, presumably due to direct photolysis. The W spectrum of DBCP shows only weak absorbance a t 254 nm ( E (base 10)= 14.7M-lcm-l), but apparently this is enough to yield some photolysis when hydrogen peroxide is at low concentrations. However, under the conditions used in this work, where [DBCP], was always 1.2 pM or less, direct photolysis cannot be a factor except when the hydrogen peroxide concentration is at a comparable level. The data in Table 2 may be used to calculate the value of the quantum yield for direct photolysis of DBCP. Although the uncertainty is large since the rates are quite slow, the calculated value of QDBCP is 0.49(f0.07). &DBCP,OH for Dibromochloropropane. It is possible to determine the value of the rate constant for the
Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 2319 120
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UV INTENSITY, W/L Figure 4. Effect of 254 nm UV intensity on the pseudo-first-order rate constants for the oxidation of DBCP (4 mM bicarbonate). [HzOzl0: 0 , i . O mM; 0 , 1.5 mM. Table 2. Direct Photolysis of DBCP W intensity (WiL) k, ( x 105 6-11 ~
0"
0.12" 0.24" 0.36" 0.48" 0 0.12 0.24 0.36 0.48
NR 6.3 12.0 17.0 27.9
NR 8.5 13.0 18.3 36.6
Ob = kd2.303cbZ0
0.45 0.43 0.40 0.50 0.61 0.46 0.44 0.65
a Runs made with no gas flow; others made with 1Umin oxygen gas flow. b = 15.8 cm; QBCP = 14.7 M-l cm-'; NR = no reaction detected. Average O = 0.49 f 0.07.
reaction of a substance with OH radicals by several methods. One may carry out a competitive kinetic experiment in which a substance with a known rate constant with the OH radical is mixed with the compound of interest, provided no reactions other than with 'OH affect either substrate. The ratio of rate constants is then obtained and the value of k M , o H calculated for the substrate of interest. In a run with [DBCP], and [PCEI, at approximately 1.2 pM, the rate constants for ~ for PCE and oxidation were 35.4 and 2.25 ( x ~ O - s-l) DBCP, respectively. Using this data, the value of kDBCp,OH derived is 1.4 x lo8 M-l s-l if the value for k p c E , o H is taken as 2.3 x lo9 M-l s-l (Buxton et al., 1988). Also, it is possible t o obtain the same parameter from the absolute value of k, obtained from the ozone1 peroxide system in bicarbonate-buffered distilled water (Glaze and Kang, 1989a,b), provided the mass transfer coefficient of the reactor is known. In this case, the
reactor used was the same one used by Glaze and Kang, and the value of kDBCp,OH obtained was 1.47 x lo8 M-l s-l. The value of kDBCp,OH used in the following model is 1.47 x lo8 M-l s-l. The value of kEDB,OH for EDB was also measured in these experiments. The value obtained by competition was 1.6 x lo8M-l s-l, very close to the value for DBCP which is structurally analogous. Derivation of a Kinetic Model for the H 2 0 W Process. From the reactions listed in Scheme 1, a kinetic model for the H202/LTV process, predicting the concentrations of DBCP and peroxide as a function of time, initial concentration of peroxide, and CT,has been derived. The model is similar to but somewhat more detailed than the model presented by DeLatt et al. (1988). In the model, two processes are presumed to be responsible for the oxidation of organic compounds: reaction with the hydroxyl radical formed by photolysis of H202 and direct photolysis of the organic by ultraviolet radiation. In the first approximation, applicable to trace levels of contaminants of poor UV absorbers, no interaction between substrate oxidation byproducts with reactants or the W field is considered. Moreover, since natural waters usually contain significant amounts of bicarbonate, the first-order model assumes that HC03- and C032- are the major scavengers of OH radicals other than substrate and H202. In this paper, we have also assumed that the carbonate radical anion, the product of the reaction of carbonate or bicarbonate with the OH radical, does not react with substrate and that its principal fate is reaction with hydrogen peroxide (see eq 2). In subsequent work, we will extend the model to cover cases where the organic compound is a
strong absorber and where carbonate radical chemistry may conceivably play a different role. With these provisos, for a monochromatic radiation oxidation is represented source, the rate of substrate (M) by eq 3
where ~ M , O Hrefers to the rate constant for reaction of M with OH radicals, assumed t o be a t a steady state concentration ['OH],,, (PM is the quantum yield for the photolysis of substrate M, f~ is the fraction of radiation absorbed by M with molar absorptivity EM (eq 41, I, is the intensity of the W radiation, and At is the total absorbance of the solution (eq 5), all at the wavelength I of the radiation source. With the assumptions stated above, fH& x 1 and
However, as noted above, the absorbance of the DBCP under conditions used in this study is always negligible compared to that of the hydrogen peroxide. Thus, the term for direct photolysis of DBCP contributes very little except when peroxide is absent. The model assumes a steady state concentration for the various intermediates ('OH, HOz' 02'- and 'cosradicals), calculated by setting the rate of radical production equal to the rate at which they are destroyed. Expressions for the steady state concentration of the three key radicals POH, 02'-and 'COa-) are shown in Scheme 2. Initial concentrations for hydrogen peroxide, bicarbonate, pH, and DBCP were assumed, and values of Io, E, KDBCP,OH, and b were taken from the measurements described above. Initial levels of all radicals were taken as zero. Using the set of equations listed in Scheme 2, values for the steady state concentrations of the radicals and values for [Hz021and [DBCP] were calculated from the model at t = 1s. These values were substituted into the same equations and the equations solved for concentration values a t t = 2 s. This procedure was repeated for 1s intervals until the total
Table 3. Comparison of Observed and Predicted Pseudo-First Order Rate Constants of DBCP Oxidation
2804 2704 2604 2004 1904 1304 2304 2404 2504 3004 3104 3204 3304 3404
0.054 0.096 0.35 0.56
1.00 1.50 3.00 4.40 6.60 0.26 0.50 0.79 1.50 3.00
0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48 0.48
4 4 4 4 4 4 4 4 4 0.1 0.1 0.1 0.1 0.1
8.4 8.3 8.4 8.5 8.4 8.4 8.4 8.4 8.5 8.1 8.4 8.0 8.2 8.2
27 31 62 82 106 108 107 86 71 258 272 271 229 160
4 8 15
34 56
70 81 74 58 236 267 274 236 172
elapsed reaction time was 1440 s. The calculated values of DBCP a t various times were then used to calculate KO values by loglinear regression. Overall, the correlation of the predicted and observed values of [H2021tr [DBCPlt, and K O for a total of 33 experimental runs under various conditions was good. Predicted pseudo-first-order rate constants of DBCP oxidation were calculated from the slope of d In [DBCPlt vs time, where DBCP concentrations were taken from the predictions of the computer model. The average deviation of these values of k, from the experimental values for 33 runs of the study was 19.7% and the range of derivations was from -39.2 to +40.5%. Table 3 shows data for runs using four UV lamps at 4 mM and 0.1 mM CT.The table compares measured and calculated values of k,. Figure 5 shows observed and predicted values of the peroxide concentration for two representative runs. It is clear from Table 3 that the model underpredicts the rate of decomposition of DBCP at 4 mM CTexcept at high concentrations of H202. The serious discrepancies between experiment and model predictions are most apparent at low peroxide levels where the rates are extremely slow and subject t o experimental errors, e.g., due t o small amounts of sparging. The dependence of the rate on initial peroxide concentration is correctly predicted by the model, however, although the optimum initial concentration of H202 is too high. At the lower CTvalue, the agreement between the model and experi-
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2321
04 0
I
400
800
.
I
1200
1600
TIME, s
Figure 5. Comparison of measured and calculated values of hydrogen peroxide concentration in 4 mM CTsolutions with 0.48 W/L UV. [H202lo = 0.054 mM; 0, observed; 0 , predicted. [Hz0~1~ = 0.096 mM; A, observed; A, predicted.
mental results is excellent. The fact that the model is inaccurate at high levels of carbonate but not at low CT may indicate that another decomposition mechanism is operating at high levels of carbonate alkalinity, perhaps due to carbonate radical ion reactions with DBCP. Effect of CTand pH. The effect of CTis illustrated in the experimental data in Tables 1and 3. The values of k, decrease with increasing CT a t a given pH. This observation is consistent with the first equation in Scheme 2 which shows that [OHIssand thus k, should be inversely proportional to CT when carbonate and bicarbonate are the principal OH scavengers. This is confirmed in Figure 6 which shows data for runs a t 1.00 mM [HzOzl, and 0.48 W/L W, all near a pH of 8. Glaze and Kang (1989a,b) have shown that this inverse proportionality with CTis also observed in the oxidation of PCE with OH radicals produced by the 03/HzOz process. Effect of pH. As noted above, inorganic carbon species have different rate constants with OH radicals. The values for carbonate and bicarbonate ions are shown in Scheme 1. The value for carbonic acid is too slow to measure. The last data entries in Table 1show that the value of k, decreases with increasing pH a t otherwise constant conditions including 4 mM CT.This is due to the fact that the carbonate ion is 49 times more effective at OH scavenging than bicarbonate (Buxton et al., 1988). Effect of Nitrate, Sulfate, Chloride, and Phosphate Anions. We have not included the effects of these anions in the model because the reaction of the corresponding radicals with DBCP is unknown. How-
ever, over a limited concentration range, the rate of DBCP oxidation is not effected by chloride, sulfate, and phosphate ions, but nitrate has a significant retarding effect (Table 4). This is apparently due to the strong absorbance of nitrate a t 254 nm which shields photons from hydrogen peroxide. This may be an important deterrent to the use of the H202/UV process in some groundwaters that are high in nitrate ion. Oxidation of DBCP in a Natural Groundwater. Table 5 shows the results of studies in which DBCP was photooxidized in a Southern California groundwater (North Hollywood well no. 27; alkalinity, 180 mg/L as CaC03 or 3.6 mM [HCOs-I; pH 7.9). It is clear from a comparison of this data with that in Table 1 that the rate of substrate oxidation is slower in the natural water than in distilled water although the bicarbonate level is somewhat lower in the groundwater. For runs under similar conditions, the rate of decomposition of DBCP is slowed by about 25% compared to that of distilled water. Apparently, other substances are present in the groundwater that affect the rate of oxidation, probably through W light shielding and OH radical scavenging. Kang (1989) has used these same waters in studies in which tetra- and trichloroethylene were oxidized with the 03/HzOz process and has estimated that substances other than bicarbonate account for approximately 30% of the OH radical scavenging. The data in Table 5 also show the accelerating effect of the softening of the water by removal of inorganic carbon by lowering the pH and purging with oxygen gas before irradiation. Summary. A kinetic model for the Hz02/1TV/DBCP/ system based on k n o w n reaction inorganic carbon (CT)
2322 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 1000
aoo
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h
cn
8
r
400
200
0 0
0.5
1
1.5
2
2.5
3
4
3.5
TOTAL CARBONATE, mM
Figure 6. Effect of total carbonate (CT) on the reciprocal of pseudo-first-order rate constants for the oxidation of DBCP with 0.48 WiL 254 nm UV radiation and 1.00 mM hydrogen peroxide. pH 8.2 f 0.2 except for run at 0.1 mM CT (pH 7.2). Table 5 . Rate Constanta of DBCP Photooxidation in NH-27O Groundwater
Table 4. Effect of Nitrate, Sulfate, Chloride, and Phosphate on Photooxidation of DBCP ~~
[HzOzl (mM)
U? intensity
[CTI (mM)
[NOa-I (mM)
2.0 2.0 2.0 2.0
2.0 5.0 10 25
1.00 1.00 1.00 1.00
(WiL) 0.48 0.48 0.48 0.48
[HzOzl (mM)
W intensity WiL)
[CT] (mM)
[SO42-]
1.00 1.00 1.00
0.48 0.48 0.48
2.0 2.0 2.0
1.0 2.0 5.0
[HzOzl (mM) 1.00 1.00 1.00
UV intensity (W/L) 0.48 0.48 0.48
[CTI (mM) 2.0 2.0 2.0
[Cl-I (mM) 1.0 2.0 5.0
1.00 1.00 1.00
0.48 0.48 0.48
2.0 2.0 2.0
1.0 2.0 5.0
(mM)
ko
(xlO5s-1) 85.4 59.5 38.2 25.5 ko (~1O~s-l)
170 169 172 ko (x105s-l)
169 164 152
143 149 166
chemistry and photochemistry agrees well with laboratory scale experimental data, especially at low inorganic carbon levels. The poorer agreement at high CTvalues suggests that another mechanism for DBCP decomposition may occur under these conditions such as reaction with carbonate ion radical. This and similar studies show that advanced oxidation processes may be modeled on the basis of a knowledge of the chemistry of the
0.0 0.10 0.50 1.0 3.2 5.6 1.0 1.0 1.0 1.Ob Lob 1.oc
0.48 0.48 0.48 0.48 0.48 0.48 0.24 0.36 0.48 0.48 0.48 0.48
3.6 3.6 3.6 3.6 3.6 3.6 3.6 3.6 1.8 0.1 0.1 2.1
7.9 7.9 7.9 7.9 7.9 7.9 7.9 7.9 6.4 5.0 5.0 8.3
20.0 19.4 41.4 87.8 53.3 32.4 29.6 56.7 141 168 191 133
North Hollywood well no. 27. pH lowered to 5.0 with HC1, 30 min purge with 1 Umin oxygen gas before oxidation. pH lowered to 5.0 with HC1,30 min purge with 1 Umin oxygen gas, and final adjustment of pH with NaOH before oxidation.
processes. These models are useful as a first approximation and may be used by practitioners to determine if an AOP is potentially useful for treatment of a specific water for specific objectives. Experimentation will be necessary to verify the model predictions, and in some cases, for example with complex wastewaters, models may be of little benefit. This work has also shown that the reaction of OH radicals with DBCP, and by analogy with other chlorinated alkanes, is a slow process. Thus, AOPs based on hydroxyl radical chemistry alone will probably not be cost-effective unless other OH scavengers such as inorganic carbon species and natural organic matter are
_+-
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2323 at low levels in the water to be treated. This suggests that presoftening of high-alkalinity groundwaters may be required in AOPs to be used for chloroalkane removal. Alternatively, more intense and lower wavelength U V lamps may be used for direct photolysis of the chloroalkanes.
Acknowledgment The experimental work for this paper was carried out at UCLA. The support of several organizations is acknowledged: the State of California Department of Health Services, the City of Los Angeles Department of Water and Power (LADWP), the U.S. Environmental Protection Agency through Cooperative Agreement CR813188 to LADWP, and the National Science Foundation through a grant to the UCLA Center for Hazardous Substances Control. This article has not been subjected to review by EPA and, therefore, does not reflect the views of this agency, and no offical endosement should be inferred. We are also grateful to the DuPont Co. for a grant to J.W.K. a t UNC/CH.
Literature Cited Baxendale, J. H.; Wilson, J. A. Photolysis of Hydrogen Peroxide at High Light Intensities. Trans. Faraday SOC. 1957,53,344356. Berglind, L.; Gjesing, E.; Skipperud Johansen, E. Removal of Organic Matter in Water by W and Hydrogen Peroxide. In Oxidation Techniques in Drinking Water Treatment; Kuhn, W., Sontheimer, H., Eds.; U.S. Environmental Protection Agency, EPA-570/9-79-020, U.S. Government Printing Oflice: Washington, DC, 1979;pp 510-523. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical Review of Data Constants for Reactions of Hydrated Electrons, Hydrogen Atoms and Hydroxyl Radicals in Aqueous Solutions. J . Phys. Chem. Ref. Data 1988,17,513-886. Draganic, Z.D.; Negron-Mendoza, A,; Sehested, K.; Vujosevic, S. I. Radiolysis of Aqueous Solutions of Ammonium Bicarbonate Over a Large Dose Range. Radiat. Phys. Chem. 1991,38,317321. Elliot, A. J.; Buxton, G. V. Temperature Dependence of the Reaction OH 0 2 and OH HO2 in Water up to 200 "C. J. Chem. SOC.,Faraday Trans 1 1992,88,2465-2470. EPRI. Proceedings of the First International EPRZI NSF Symposium on Advanced Oxidation; EPRI TR-102927,Electric Power Research Institute, Palo Alto, CA, 1994;Vol. 2. Eriksen, T. E.; Lind, J.; Merenyi, G. On the Acid-Base Equilibrium of the Carbonate Radical. Radiat. Phys. Chem. 1985,26,197199. Glaze, W. H.; Kang, J. W. Advanced Oxidation Processes. Description of a Kinetic Model for the Oxidation of Hazardous Materials in Aqueous Media with Ozone and Hydrogen Peroxide in a Semibatch Reactor. Znd. Eng. Chem. Res. 1989a,28, 15731580. Glaze, W. H.; Kang, J. W. Advanced Oxidation Processes. Test of a Kinetic Model for the Oxidation of Organic Compounds with
+
+
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Received for review August 15, 1994 Accepted January 13,1995@ IE9404902 Abstract published in Advance ACS Abstracts, April 1, 1995. @
