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Ind. Eng. Chem. Res. 1989,28, 1580-1587
Stand. Ref. Data Ser. (U.S. Natl. Bur. Stand.) 1977, No. 59. Forni, L.; Bahnemann, D.; Hart, E. J. Mechanism of the Hydroxide Ion Initiated Decomposition of Ozone in Aqueous Solution. J . Phys. Chem. 1982,86,255. Glaze, W. H. Drinking-Water Treatment With Ozone. Enuiron. Sci. Technol. 1987,21,224-230. Glaze, W. H.; Kang, J. W. Advanced Oxidation Processes for Treating Groundwater Contaminated With TCE and P C E Laboratory Studies. J . A WWA 1988,80,57-63. Glaze, W. H.; Kang, J. W. Advanced Oxidation Processes. Test of a Kinetic Model for the Oxidation of Organic Compounds with Ozone and Hydrogen Peroxide in a Semi-batch Reactor. Znd. Eng. Chem. Res. 1989,following paper in this issue. Glaze, W. H.; Kang, J. W.; Aieta, E. M. Ozone-Hydrogen Peroxide Systems for Control of Organics in Municipal Water Supplies. In The Role of Ozone in Water and Wastewater Treatment, Proceedings of the Second International Conference; TekTran International, Ltd.: Edmonton, Alberta; p 233 April 28-29,1987a. Glaze, W. H.; Kang, J. W.; Chapin, D. H. The Chemistry of Water Treatment Processes Involving Ozone, Hydrogen Peroxide and Ultraviolet Radiation. Ozone Sci. Eng. 1987b,9,335. Glaze, W. H.; Koga, M.; Cancilla D. Ozonation Byproducts. 2. Formation of Formaldehyde and Other Carbonyl Compounds by Ozonation of Drinking Water. Enuiron. Sci. Technol. 1989,in press. Gurol, M. D.; Nekouinaini S. Effect of organic substances on mass transfer in bubble aeration. J. WPCF 1985,57(3),235-240. Haber, F.; Weiss, J. The Catalytic Decomposition of Hydrogen Peroxide by Iron Salts. Proc. R. SOC. A 1934,147, 332. Holcman, J.; Sehested, K.; Hart, E. J. Rate Constants and Product of the Reactions of e, O;, H and OH with Ozone in Aqueous
Solutions. Radiat. Res. Proc. Znt. Congr., 7th 1983,Paper A2-12. Koester, R.; Asmus, K. D. Die Reaktionen chlorierter Aethelene mit hydratisierten Elektronen und OH-Radikalen in waessriger Loesung. 2. Naturforsch., B. 1971,26(11), 1108. Kosak-Channing, L. F.; Helz, G. R. Solubility of Ozone in Aqueous Solutions of 0-0.6 M Ionic Strength at 5-30 "C. Enuiron. Sci. Technol. 1983,17, 145-149. Lazrus, A. L.; Kok, G. L.; Gitlin, S. N.; Lind, J. A.; Mclaren, S. E. Automated Fluorometric Method for Hydrogen Peroxide in Atmospheric Precipitation. Anal. Chem. 1985,57,917-922. Nakayama, S.; Esaki, K.; Namba, K.; Taniguchi, Y.; Tabata, N. Improved Ozonation in Aqueous Systems. Ozone: Sci. Eng. 1979, I, 119. Peyton, G. R.; Glaze, W. H. Mechanism of Photolytic Ozonation. In Photochemistry of Environmental Aquatic Systems; Zika, R. G., Cooper, W. J., Eds.; ACS Symposium Series 327, American Chemical Society: Washington, DC, 1986; pp 76-88. Staehelin, J.; Buhler, R. E.; Hoign6, J. Ozone Decomposition in Water Studied by Pulse Radiolysis. 2. OH and H 0 4 as Chain Intermediates. J. Phys. Chem. 1984,88,5999-6004. Staehelin, J.; HoignB, J. Decomposition of Ozone in Water: Rate of Initiation by Hydroxide Ion and Hydrogen Peroxide. Enuiron. Sci. Technol. 1982,16, 676. Taube, H. Photochemical Reactions of Ozone in Solution. Trans. Faraday SOC. 1957,53,656. Weeks, J. L.; Rabani, J. The pulse radiolysis of deaerated aqueous carbonate solutions. I. Transient optical spectrum and mechanism. 11. pK for OH radicals. J. Phys. Chem. 1966,70(7),2100. Received for review January 31, 1989 Accepted June 28, 1989
Advanced Oxidation Processes. Test of a Kinetic Model for the Oxidation of Organic Compounds with Ozone and Hydrogen Peroxide in a Semibatch Reactor William H. Glaze*?'and Joon-Wun Kangt Environmental Science & Engineering Program, School of Public Health, University of California, Los Angeles, California 90024
Experimental data are presented to test a kinetic model of the 03/H202process in a semibatch reactor. The effect of bicarbonate and carbonate ions is measured and found to be in concurrence with model predictions. The effect of p H in the ozone mass-transfer-limited region was examined in bicarbonate-spiked distilled water. Since the reaction is mass transfer limited, the primary effect above pH 7 is the result of changes in the distribution of inorganic carbon species which are OH-radical scavengers. Below p H 7, there is a lag period during which ozone and peroxide increase until the chain reaction begins. The effects of chloride ion and the concentration of radical scavengers other than carbonate species in ground waters are also measured. The mass-transfer/reaction rate model has been used to estimate rate constants for the reaction of hydroxyl radicals with trichloroethylene, carbon tetrachloride, and two bicyclic alcohols, 1,2-dibromoethane, 1,2-dibromo-3-chloropropane, 2-methylisoborneol and geosmin. In the preceding paper (Glaze and Kang, 1989),the rate constants derived from fundamental kinetic studies along with mass-transfer principles were used to derive a kinetic model for the oxidation of organic substances by the 03/H202process in a semibatch reactor. The principal species and equations of the model are shown in Schemes I and 11, respectively. In brief, the 03/H202system consists of a rate-determining reaction between ozone and the conjugate base of hydrogen peroxide. A set of rapid re-
* To whom correspondence should be addressed. t Current address: Department of Environmental Sciences and Engineering, School of Public Health, The University of North Carolina at Chapel Hill, Chapel Hill, NC 27599-7400.
0888-5885/89/2628-1580$01.50/0
actions takes place thereafter, yielding the principal reactive intermediate, the hydroxyl radical. The OH radical may oxidize substrate M, but because of its high reactivity, it also may react with a variety of other substances commonly found in natural water. In the model developed in the previous paper, the principal OH scavengers were assumed to be the bicarbonate and carbonate ions, with hydrogen peroxide (or its conjugate base) and ozone playing more or less secondary roles depending upon their concentrations in the reaction medium. In a semibatch reactor, the model predicts that the rate of substrate (M) oxidation at a given ozone dose rate will increase linearly with an increasing rate of addition of peroxide (referred to as region I), that the oxidation rate 0 1989 American Chemical Society
Ind. Eng. Chem. Res., Vol. 28, No. 11,1989 1581 Scheme I. Diagram of the Principal Reactions and Species Involved in the OI/HzOz Process in Distilled Water
I
H A Wi (11
?
of M will level off after the Hz02/03ratio passes through its stoichiometricoptimum, after which the system is Ozone transfer limited (region 111),and that the mass-transferlimited rate in region I11 will be proportional to the concentration of ozone in the gas phase and the mass-transfer coefficient of the reactor and inversely proportional to the sum of the kinetic factors for all OH-radical scavengers. While the model developed for the distilled water system was successful in predicting the rate of tetrachloroethylene (PCE) oxidation and the concentration of residual ozone and peroxide in regions I and 111, respectively, there are several features of the model that remain unresolved when the matrix is changed to a real surface or ground water. This and subsequent papers will investigate these effects.
Experimental Section Methods. The experimental procedures used in this work are the same as those described in the previous paper (Glaze and Kang, 1989) except as follows. The effect of bicarbonate was studied with a series of measurements in distilled water (Arrowhead Water Company, Monterey Park, CA) into which were spiked concentrated solutions of sodium bicarbonate (Fisher, ACS certified grade) and the organic compound of interest so as to achieve the target concentration of each. The following organic compounds were used with the noted target concentrations: PCE,
tetrachloroethylene (Aldrich, laboratory grade; 50-70 pg/L); trichloroethylene (Fisher Scientific; 500 pg/L); DBCP, 1,2-dibromo-3-chloropropane(Columbia Organic Chemicals, laboratory grade; 200 pg/L); EDB, 1,2-dibromoethane (Eastman Kodak, laboratory grade; 200 pg/L); DBP, 1,2-dibromopropane (Eastman Kodak, laboratory grade; 300 pg/L in hexane); CTC, carbon tetrachloride (Aldrich, HPLC grade; 200 pg/L); geosmin (trans-l,lO-dimethyl-trans-9-decalol) and 2-methylisoborneol (Wako Chemicals; 200 ng/L). For some runs, the pH of the test water sample was adjusted prior to oxidation with either a 1 N HC1 or 2 N NaOH solution. Analytical Methods. Tetrachloroethylene, trichloroethylene, 1,2-dibromo-3-chloropropane, and carbon tetrachloride were analyzed by microscale liquid-liquid extraction as described in the previous paper (Glaze and Kang, 1989). Geosmin and 2-methylisoborneol were analyzed by the closed-loop stripping analysis method (Krasner et al., 1981; Glaze et al., 1989). In this technique, volatile organic compounds of intermediate molecular weight were stripped from the water by a recirculating stream of air. The stripped TAO compounds were trapped on approximately 1.5 mg of granular activated carbon (GAC). The carbon was eluted with approximately 25 pL of methylene chloride. The extract was analyzed by GC/MSD (Hewlett-Packard Model 5890/5970), which utilized a DB-5 fused-silica column (J & W Scientific). The initial temperature was 33 "C (1 min), to 92 "C at 4 deg/min and then to 230 "C at 10 deg/min. Materials. Three different sources of water were used during these studies: distilled water (Arrowhead Water Co., Monterey Park, CA) and two ground waters taken from wells No. 14 and 26 of the North Hollywood field of the Los Angeles Department of Water and Power (DWP). Ground water was collected by DWP personnel in 55-gal polyethylene-lined drums. For most runs, tetrachloroethylene (PCE) was spiked into the water to raise its concentrations to about 50-70 pg/L. Typical water quality parameter values of North Hollywood wells No. 14 and 26 are shown in Table I.
Results and Discussion Effect of Natural Water Contaminants on Oxidation Processes. Three generic types of substances may be present in natural waters that will affect the 03/H202 process: initiators and promoters of the ozone decompo-
Scheme 11. Summary of Kinetic Equations for OI/HIOz Process Model in Bicarbonate-SpikedDistilled Water
region I
where region I11
where region I1
1582 Ind. Eng. Chem. Res., Vol. 28, No. 11, 1989 Table I. Test Water Quality Data' alkalinity, mg as CaC03/L hardness Ca, mg/L Mg, mg/L Fe, mg/L C1-, mg/L turbidity PO4, mg/L BOD TOC PH
NH-14 300 422 121 29 KO.01 21 0.1
0.08 0.5 0.5-0.7 1.2-7.6
NH-26 200 420 108 37 ND 37 0.05 0.05 0.6 0.4 7.6-7.8
Water quality data provided by the City of Los Angeles Department of Water and Power.
-35-
A
-4 3
-
C
5
10
'5
20
25
- min Figure 1. Effect of bicarbonate on the rate of oxidation of PCE by the 03/H202process in distilled water. Ozone dose, 0.23 mg/(L min); peroxide dose, 0.20-0.24 mg/(L min). Time
Table 11. Rate Constants of Some Inorganic and Organic Compound with OH Radicals" compd M kM,oH, M-' s - ~ HC031.5 x 107 c0:4.2 X lo8 cu2+ 3.6 X lo8 Fez+ 1.2 x 109 Mn2+ 1.4 X IO8 I1.0 x 1010 Br109-1010
c1-
io i
-1.0
300
-1.5
1 0 0 ~
-2.0 0
2
1
3
4
5
6
[HC03-], mM
-2.5 a
Figure 2. Reciprocal of ko versus bicarbonate concentration. Ozone dose, 0.23 mg/(L min); peroxide dose, 0.20-0.24 mg/(L min); pH 7.6-7.8. Table IV. Effect of Chloride Ion on the Rate of Oxidation of PCE" NaCl. me/L k,. lo4 14.1 0 13.8 50 15.4 100 15.1 500 13.3 5000
"08 dose: 0.23 mg/(L min). H202dose: 0.09 mg/(L min). pH 8.4-8.5. Total carbonate: 4 mM.
Figure 2 is a plot of the data in Table 111, the slope of which should be equivalent to the term (k13 + k14(lopH-pX'))/([kM,OH][kLu](P/H)) in eq 8. The linear regression equation for Figure 2 is 1O3(l/ko) = 94.1[HC0
