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ozone–peroxide oxidation using FT-IR
PII: S0043-1354(00)00162-7
Wat. Res. Vol. 34, No. 16, pp. 4036±4048, 2000 7 2000 Elsevier Science Ltd. All rights reserved Printed in Great Britain 0043-1354/00/$ - see front matter
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MINERALIZATION AS A MECHANISM FOR TOC REMOVAL: STUDY OF OZONE/OZONE±PEROXIDE OXIDATION USING FT-IR STEVE A. CARR* and RODGER B. BAIRD County Sanitation Districts of Los Angeles County, San Jose Creek Water Quality Laboratory, 1965 S. Workman Mill Road, Whittier, CA 90601-1415, USA (First received 1 July 1999; accepted in revised form 7 January 2000) AbstractÐFT-IR spectroscopy was used to study the course of abiotic mineralization in samples treated with ozone and ozone:peroxide. The importance of oxidative mineralization as a mechanism for dissolved organic carbon (DOC) removal during disinfection must be determined if advanced oxidation processes (AOPs) in water and wastewater treatment is to be fully exploited. Understanding the full impact of these applications is critical for determining how eective a remedy AOPs are for eliminating disinfection by-product formation and for assessing how well these applications might facilitate compliance with future regulatory changes. Although these processes have been shown to surpass the level of disinfection aorded by traditional techniques and improve biodegradability, this study found that the use of ozone, or ozone in combination with hydrogen peroxide, to chemically eect TOC reduction may have practical limitations when applied to dilute aqueous solutions containing both refractory and degradable dissolved organic carbons. Reactions possessing very favorable thermodynamics were found, by an indirect measure, to be complicated by solute/solvent interactions which may have served to inhibit oxidation. It was also shown that ``close to real time'' CO2 plots generated using highly sensitive FT-IR techniques may have some practical utility in assessing mineralization potential and optimizing disinfection processes. 7 2000 Elsevier Science Ltd. All rights reserved Key wordsÐTOC, ozonolysis, mineralization, b-cyclodextrin, infrared, advanced oxidation, hydroxyl radical, disinfection, solvation eect
INTRODUCTION
When conventional water treatment techniques are employed, dissolved organic matter can be removed from water through a combination of biological, physical and chemical means (Weber, 1972; Jolley et al., 1978; Brunet et al., 1982). These processes are known to interact and exhibit complex, changing synergies as treatment progresses. Of the three mechanisms, biodegradation is generally acknowledged to be the primary removal process, for both natural organic matter (NOM) and euent organic matter (EfOM) in treated wastewaters. This is followed by physical adsorption processes which include granular activated carbon, coagulation and settling. Chemical oxidation, by comparison, is believed to make only minor, direct contributions to dissolved organic carbon (DOC) removal, when traditional disinfection techniques are utilized. Ozone, the most powerful of the common disinfectants, has been reported to be very eective at *Author to whom all correspondence should be addressed.
reducing disinfection by-product (DBP) formation (Miltner et al., 1992; Anderson et al., 1990), in addition to enhancing ¯occulation and oxidizing taste and odor compounds in water (Ferguson et al.,1990). In Europe, where it has been used successfully as a disinfectant for several decades, experience has shown that ozonation can alter TOC levels, and produce large concentrations of lowmolecular-weight carbon compounds (Kuo et al., 1978; Legube et al., 1980; Glaze et al., 1987; Hacker et al., 1994). More recently, advanced oxidation processes (AOPs) which utilized ozone as one of the components, have been shown to increase the destruction of complex organic constituents which are refractory to other treatment (van der Kooij et al., 1989; Ferguson et al., 1990; Brunet et al., 1982). Partially oxidized by-products of these AOP reactions have also been shown to exhibit greater biodegradability (Amy et al., 1996; Duguet et al., 1985; Narkis and Schneider-Rotel, 1980; Legube et al., 1980) and have higher adsorption coecients (Janssens et al., 1984), probably due to their increased polarities. Although oxidative mineralization is sus-
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FT-IR study of DOC mineralization
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Fig. 1. Ozonolysis apparatus.
pected to be a factor in the observed reduction of DBP and TOC in treated waters (Xu et al., 1989; Adams et al., 1997; Volk et al., 1997; Caprio et al., 1987), there have been few investigations to selectively study its mechanisms and its total contribution to TOC removal. This paper will evaluate the ability of ozone and ozone±peroxide to chemically remove dissolved, non-purgeable organics in water through direct oxidative mineralization. MATERIALS AND METHODS
Experiments were carried out at room temperature in a batch mode. Bench scale ozonation was performed using a 2.2-l Pyrex column (1.07 m 54 cm) ®tted with a sintered glass inlet diuser (Fig. 1). The reactor column was constructed of glass; all other ®ttings and tubing were stainless steel or Te¯on. Flow rates to the column were controlled using an in-line pressure regulator in series with adjustable needle valves. The ¯ow rate was measured using a J&W digital ¯ow meter. The ¯ow rate of the ozonized gas from the generator was adjusted to 433 ml/min. This rate permitted a minimum of three volume exchanges in the IR gas cell and one complete displacement of the apparatus dead volume1 per infrared spectra. Infrared spectra were obtained on reaction products using a Mattson Galaxy 5000 FT-IR spectrophotometer equipped with a DTGS detector. A cylindrical glass Flow-Thru Gas Cell (110 30 mm) with Te¯on stopcock, Viton O-ring seals and two CaF2 IR-Transmission Windows (38 6 mm), having a transmission range from 1100 to 50,000 cmÿ1 (9.0±0.20 mm) was used to analyze the gas-
1
(apparatus dead volume = column head volume+tubing volume+gas cell volume).
eous reaction euent. CaF2 windows were selected because of their resistance to water, acids and alkalies. These properties are important because the cell windows are exposed to a continuous stream of moist air which may contain traces of oxidized by-products. Windows were frequently polished between runs, using aluminum oxide powder (<50 mm) and water to remove any surface build-up. Infrared spectra were acquired using the following parameters: starting wavenumber 1350 cmÿ1; ending wavenumber 4000 cmÿ1; resolution of 4.0 cmÿ1; number of scans 64; gain 1; forward velocity 10.0 KHz; reverse velocity 10.0 KHz; time per spectra 1.23 min; apodization usedÐtriangulation. All compounds used in this study were reagent grade or better, and were purchased from Aldrich, Baker or Mallinckrodt. The b-cyclodextrin (b-CD) was donated by Cerestar, Inc. Standards were prepared in a 20-l pyrex carboy at 10 mg/l concentrations, except when solubility was a limiting factor, using Barnstead Nanopure Type II water. After preparation, standards were aerated to aid dissolution, provide mixing and remove residual bicarbonate. In general standards were not acidi®ed or buered prior to oxidation (except where speci®ed). Well waters, groundwater and plant euent samples were acidi®ed to pH R 3 before aeration. Portions of the standard not used on the day of preparation were capped and refrigerated at 48C. Prior to the start of a run, samples were purged completely of carbon dioxide using CO2-free compressed air from a Balston Model 7562 FTIR Purge Gas Generator. The ozonation apparatus was cleaned by ¯ushing with DI water and rinsing, with the standard to be tested, for a minimum of 5 min. After cleaning, the column was ®lled with sample using a FMI pump. A continuous CO2-free air purge was maintained through the apparatus, before and after ®lling, to remove residual CO2 and prevent ambient contamination. For this study, ozone was generated from air or pure oxygen using a PCI Ozone Generator (Model GL-1, West Caldwell New Jersey). The ozone generator was turned on
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Steve A. Carr and Rodger B. Baird
for a minimum of 15 min prior to the start of a run to allow the output to stabilize. Ozone concentration in the feedgas [ave. concentration with O2349 mg/l, air 330 mg/ l] and the o-gas streams were measured using a PCI ozone monitor (Model LC-400, West Caldwell, NJ). The ozonized gas stream from the generator was fed continuously through the bottom of the reactor as micro-bubbles for the duration of a run. Sample mixing was accomplished solely via this process. Earlier dye studies determined that the turbulence and agitation resulting from aeration of the reactor column provided adequate mixing. A DI water blank was determined to have a dissolved ozone concentration between 5 and 8 mg/l with an average value of 07 mg/l, using the indigo blue method. The dis-
solved ozone concentration of a plant euent sample was 5 mg/l after 5.0 min contact time. The euent gas stream from the reaction column was conveyed via a 1/40 Te¯on tubing to the gas ¯ow cell in the FT-IR sample chamber. Blank background spectra of the Gas Cell were collected using CO2-free air from the Balston generator. The o gas from the column was then monitored by infrared for CO2 depletion. A column background spectrum was collected when the CO2 infrared absorbance band disappeared. This reference was stored in instrument memory just prior to the addition of the O3 oxidant, and subtracted from all subsequent spectra collected during oxidation. For ozone:peroxide experiments, 3.64 ml of a 3000 mg/l hydrogen peroxide standard was
Fig. 2. (a) Sample IR spectrum from study. (b) Sample IR spectrum with expanded region.
FT-IR study of DOC mineralization added via syringe at the start of the run, or continuously for the duration of a run using an Altec Model 110A HPLC pump connected to the column sample inlet by 1/80 Te¯on tubing (see Fig. 1.). The total concentration of hydrogen peroxide added was 5.0 mg/l. To begin an analysis, the gaseous air purge to the contact column was switched to the ozonated gas stream from the generator using a four-way stainless steel ball valve. The FT-IR was started simultaneously, initiating data collection. The FT-IR spectrophotometer was programmed to continuously collect and store multiple, incremented infrared spectra, for the duration of a run (see Fig. 2.). In addition, the Macro allowed the integrated areas for the CO2 infrared absorption bands to be stored in preassigned output ®les. For this study, 150 infrared spectra were collected over a 3 12 h period to measure CO2 concentrations, and to evaluate other constituents in the ozonated air stream after column contact. Carbon dioxide concentration was calibrated using a Scotty III Certi®ed Standard having a CO2 concentration of 408 mg/l22%, in nitrogen. A detection limit of R0.25 ppb was calculated for CO2 using serial dilutions of this certi®ed standard. The absorbance area between 2410.6 and 2281.4 cmÿ1 was used to calculate CO2 concentration (mg/l). These values were then plotted vs time using Microsoft Excel 5.0. The total evolved CO2 for a given run was determined from the integrated plots. No adjustment was made for evaporative losses which were estimated to be 32% during runs. A NaHCO3 standard spiked into an acidi®ed DI water column blank yielded >78% theoretical CO2 recovery. The b-cyclodextrin inclusion complexes with caeine, methylene blue and methyl orange were prepared in a 40 ml vial by adding 1.0 mmol of the respective reagents to 2.0 mmol slurry of b-cyclodextrin in 20 ml of DI water. The mixtures were stirred overnight prior to use. Bleaching experiments were conducted on highly colored methylene blue and methyl orange standards and their inclusion complexes. The times from the beginning of ozonation to complete color loss for the neat and complexed chromophores were compared. Two hundred and ®fty milliliters of water samples were collected in precleaned amber bottles before and after each analysis. Samples were treated with 0.5 g of sodium arsenite (NaAsO2) to remove oxyresidual, then preserved with 2 ml 1:1 H2SO4 and refrigerated. TOC analyses were conducted per Standard Methods' 19th Edition (APHA, AWWA, WEF, 1995) 5310C Heated Persulfate Oxidation Method. The analyses were run on an OI Analytical Instrument Model 1010 equipped with an autosampler OI Analytical Instrument Model 1051. RESULTS AND DISCUSSION
Bench scale experiments were performed on prepared standards, well waters from a local groundwater recharge basin, and secondary and tertiary euent samples from a wastewater reclamation plant. Conditions used in this study approximated an open, batch treatment process. Experiments were conducted with the goal of optimizing contact times to determine the capacity of ozone and ozone:peroxide to oxidatively remove dissolved organics. Knowing what occurs when DOCs are exposed to a stoichiometric excess of ozone or ozone:peroxide over prolonged periods could better de®ne the limits of advanced oxidation processes, and provide useful insights into how mineralization progresses with time.
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The study's design approximated pseudo ®rst order kinetics and optimized oxidation by keeping the ozone concentration in excess for the duration of each test. These close-to-ideal conditions were assisted by a continuous purge which forced the reaction's equilibrium to the right and removed bicarbonate ions, a potential hydroxyl radical scavenger, from the matrix. Compounds were chosen to represent the wide range of structure and functionalities which might reasonably be expected in natural and disinfected wastewater samples. Standards for this study were prepared at 10 mg/l, except when solubility was limiting. Initial concentrations for TOC were in the 1±7 mg/l range. The percent carbon mineralized was based on the molarity of carbon in standards relative to the number of moles of CO2 generated during oxidation. The experimental results for ozone and ozone:peroxide oxidation of 28 model compounds are summarized in Table 1 and Figs 3 and 4. Total CO2 concentration for these experiments was determined from integrated plots. These graphs were also evaluated to determine optimum contact times and to study relative mineralization pro®les. No measurable concentration or discernable pro®le changes were observed for CO2 plots when the feed-gas was switched between air and oxygen, even though oxygen provided >60% increase in ozone generation vs air. This suggests that the applied ozone concentration, for these experiments, was not rate limiting at or above the 30 mg/l air-feed levels. It also suggests that once initial solubility limits were attained, ozone mass transfer rates were determined solely by reaction rates, not feed-gas stoichiometry. FTIR quanti®cation of CO2 demonstrated that mineralization was not the primary or even an important mechanism for TOC removal from standards. Even after unrealistically long (03.5 h) contact with the oxidants, only 3% of the DOC, on average, was mineralized when measured as moles of CO2 generated. As can be seen in Fig. 3, neither oxidant achieved greater than 06% mineralization. The removal of TOC was about an order of magnitude higher with an average of 31%, based on the TOC concentration in standards before the application of the oxidant. Some loss of organic constituents is known to occur when water or wastewater utilities employ aeration treatments (Bell et al., 1988; Mackay and Leinonen, 1975; Smith et al., 1980; Peng et al., 1994, 1995; Roberts et al., 1983, Langlais et al., 1991). Random comparison of standards before and after air purging 030 min without ozone or ozone:peroxide, showed little or no measurable TOC changes. It is therefore unlikely that the TOC decreases observed were due to any signi®cant air stripping of the unoxidized starting materials. There was also no turbidimetric or particulate organic carbon (POC) formation indicated. The TOC data, then, were somewhat surpris-
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Steve A. Carr and Rodger B. Baird
ing, absent physical absorptive or biological processes to account for the carbon loss. TOC reduction might also be observed if partially oxidized intermediates from disinfection were resistant to oxidation under the conditions employed for TOC analysis. However, this appears unlikely. When prepared standards of low molecular weight organic acids and other known ozonation end-products were analyzed for TOC, recoveries were close
to quantitative. A more plausible explanation for the non-mineralized TOC losses may be the direct purging of oxidized intermediates from the reaction column. The persistence of multiple IR bands in the 1653±1750 cmÿ1 region (Fig. 2b) suggests the presence of aldehydes and other carbonyl bearing constituents in the post reaction stream. These products could not be resolved, however, because known by-products such as methyl glyoxal and or-
Table 1. Model compounds studied using ozone/ozone:peroxide oxidation Compound/oxidation method Acetic acid/O3 Acetic acid/OPa Adaman./O3 Adaman./OP t-Butanol/O3 t-Butanol/OP Caeine/O3 Caeine/OP t-CDTA/O3 t-CDTA/OP DEGa/O3 DEG/OP 3,5-DHBA/O3 3,5-DHBA/OP EDTA/O3 EDTA/OP Formic A./O3 Formic A./OP Glucose/O3 Glucose/OP Glycerol/O3 Glycerol/OP Humic A./O3 Humic A./OP Inositol/O3 Inositol/OP KHP/O3 KHP/OP 1-NAM.HCl/O3 1-NAM.HCl/OP 1-NED/O3 1-NED/OP Oxalic A./O3 Oxalic A./OP PDT/O3 PDT/OP PEG/O3 PEG/OP PCP/O3 PCP/OP Phenol/O3 Phenol/OP PMA/O3 PMA/OP Quinaldic A/O3 Quinaldic A/OP Resorcinol/O3 Resorcinol/OP Tannic A./O3 Tannic A./OP Tartaric A./O3 Tartaric A./OP TBAC/O3 TBAC/OP TPM/O3 TPM/OP a
Init. [C] mmol.
[CO2] mmol.
Init. TOC (mg/l)
Final TOC (mg/l)
[TOC]/mmol C
D [C] %
D [TOC] %
Qrfx
0.765 0.765 1.453 1.453 1.208 1.208 0.913 0.913 0.505 0.505 0.861 0.861 1.015 1.015 0.589 0.589 0.456 0.456 0.731 0.731 0.816 0.816 0.681 0.681 0.821 0.821 0.876 0.876 1.222 1.222 1.023 1.023 0.510 0.510 0.899 0.899 0.710 0.710 0.207 0.177 1.432 1.432 1.394 1.394 1.270 1.270 1.221 1.221 0.929 0.929 0.592 0.592 1.339 0.840 1.566 1.566
0.005 0.012 0.002 0.004 0.005 0.016 0.030 0.026 0.025 0.027 0.017 0.036 0.035 0.035 0.018 0.019 0.020 0.017 0.021 0.019 0.030 0.029 0.026 0.019 0.031 0.030 0.022 0.022 0.025 0.031 0.022 0.026 0.016 0.017 0.026 0.026 0.008 0.026 0.007 0.008 0.045 0.045 0.031 0.032 0.033 0.033 0.041 0.039 0.024 0.025 0.019 0.019 0.001 0.015 0.001 0.002
4.18 4.18 7.50 7.50 1.70 1.70 3.78 3.78 4.07 4.07 4.34 4.34 5.23 5.23 3.00 3.00 2.80 2.80 3.10 3.10 4.00 4.00 2.20 2.20 4.26 4.26 3.70 3.70 6.50 6.50 4.83 4.83 2.20 2.20 4.30 4.30 4.73 4.73 0.98 0.98 7.43 7.43 6.51 6.51 6.01 6.01 6.30 6.30 4.48 4.48 3.29 3.29 6.69 6.69 6.51 6.51
3.62 2.83 9.70 7.50 2.01 2.01 0.58 0.64 0.22 0.26 1.89 0.00 0.27 0.09 0.20 0.20 0.11 0.12 0.81 0.25 0.18 0.08 0.49 0.65 0.34 0.48 1.00 0.17 2.40 1.50 1.25 1.41 0.17 0.14 0.74 0.81 4.13 1.86 0.59 0.15 0.79 0.49 1.71 1.65 1.27 0.41 3.40 3.00 1.40 1.80 0.76 0.37 6.52 4.05 1.71 0.17
5.47 5.47 5.16 5.16 1.41 1.41 4.14 4.14 8.06 8.06 5.04 5.04 5.15 5.15 5.09 5.09 6.14 6.14 4.24 4.24 4.90 4.90 3.23 3.23 5.19 5.19 4.22 4.22 5.32 5.32 4.72 4.72 4.31 4.31 4.78 4.78 6.66 6.66 4.76 5.55 5.19 5.19 4.67 4.67 4.73 4.73 5.16 5.16 4.82 4.82 5.56 5.56 4.99 7.97 4.16 4.16
0.70 1.56 0.17 0.27 0.40 1.32 3.27 2.84 5.03 5.32 1.99 4.16 3.44 3.46 3.04 3.23 4.30 3.77 2.89 2.60 3.69 3.52 3.76 2.78 3.83 3.62 2.51 2.55 2.03 2.57 2.18 2.50 3.04 3.39 2.88 2.87 1.14 3.63 3.29 4.68 3.15 3.17 2.20 2.27 2.58 2.58 3.33 3.22 2.56 2.69 3.12 3.19 0.07 1.76 0.08 0.10
13.40 32.30 ÿ29.33 0.00 ÿ18.24 ÿ18.24 84.79 83.02 94.59 93.61 56.45 100.00 94.91 98.34 93.40 93.27 96.07 95.71 73.87 91.94 95.43 97.93 77.73 70.45 92.07 88.78 72.97 95.41 63.08 76.92 74.12 70.81 92.27 93.64 82.88 81.26 12.69 60.68 39.84 84.76 89.42 93.38 73.73 74.65 78.87 93.24 46.03 52.38 68.75 59.82 76.93 88.85 2.54 39.46 73.73 97.47
1.02 0.87 1.15 1.14 1.30 0.98 0.87 0.18 0.29 0.24 1.21 1.67 0.25 0.22 0.04 0.12 0.28 0.05 0.63 0.41 0.25 0.17 0.18 0.25 0.43 0.35 0.19 0.15 0.31 0.30 0.20 0.25 0.87 0.08 0.17 0.13 1.18 1.29 0.12 0.05 0.34 0.34 0.10 0.12 0.27 0.23 0.29 0.26 0.27 0.44 0.55 0.35 1.20 0.81 0.37 0.20
Adaman=1-adamantanamine, t-CDTA=t-cyclohexylenediamine-tetraacetic acid, DEG=diethylene glycol, 3,5-DHBA=3,5-dihydroxybenzoic acid, PDT=3-(2-pyridyl)-5,6-diphenyl,-1,2,4 triazine p,p'-disulfonic acid monosodium salt, KHP=potassium hydrogen phthalate, NAM.HCl=1-naphthylamine.HCl, NED=1-naphthylethylenediamine, PEG=polyethylene glycol (MW 8000), PCP=pentachlorophenol, PMA=1-pyrenemethylamine, TBAC=tetrabutylammonium chloride, TPM=triphenylmethanol, OP=ozone:peroxide.
FT-IR study of DOC mineralization
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Fig. 3. Percent carbon mineralized for selected compounds during ozone/ozone:peroxide application.
ganic acids all possess fundamental infrared vibrational frequencies in this region (Atkinson et al., 1997). A white ¯aky build-up on the apparatus tubing was determined to be paraformaldehyde when a
sample of the material was heated above 1408C in an IR gas cell and the resulting infrared spectral bands matched those of formaldehyde. This identi®cation was later con®rmed by dissolving a portion
Fig. 4. Percent TOC change for selected compoundsd during ozone/ozone:peroxide application.
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Steve A. Carr and Rodger B. Baird
Fig. 5. The ozone/ozone:peroxide mineralization pro®les of selected model compounds.
of the heated vapors in DI water, derivatizing the solution with 2,4-dinitrophenyl hydrazine, then analyzing the hydrazone derivative by HPLC. These analyses showed unequivocally that loss of formaldehyde was occurring during ozonation, and suggest that formaldehyde loss may have been an important mode for organic carbon removal during oxidation. Similarly, the loss of CO as a by-product of DOC oxidation cannot be ruled out in light of results from previous ozonolysis studies (Niki et al.,
1987; Atkinson and Carter, 1984; Atkinson et al., 1997; Horie and Moortgat, 1991). Quantitation of the CO infrared bands at 2025 and 2240 cmÿ1 was subject to interference from an adjoining O3 band at 2097 cmÿ1 and a N2O band at 2236 cmÿ1 in the IR spectra (Kogelschatz and Baessler, 1987) (Fig. 2a). Mineralization pro®les of ozone and ozone:peroxide The data summarized in Table 1 represent the net
FT-IR study of DOC mineralization
results from 3 12 h long experiments. Although this time may be long relative to real-world disinfection processes, the duration of these experiments provided a unique window to study how AOPs progress over time. Use of an in-line FTIR also allowed the production of CO2 to be evaluated in ``close to real time'' for a range of model compounds during ozone and ozone:peroxide applications. The N2O and O3 concentration in the vented gas was monitored simultaneously as illus-
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trated in Fig. 2a. As expected, nitrous oxide was absent when oxygen was used as a feed-gas for ozone generation. Only tentative spectral assignments were made for other functional groups such as C=O, C-H, and -COOH, because of the potential for band overlap. (see IR spectra Fig. 2b). In spite of the minor role CO2 elimination had in the reduction of TOC, the variable features of the CO2 evolution plots have provided revealing insights into the complex nature of these oxidation
Fig. 5 (continued).
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Steve A. Carr and Rodger B. Baird
processes (Fig. 5a±d). Figure 5a shows the CO2 generation pro®les for selected model compounds, a tertiary wastewater euent, and groundwater samples. Figure 5b±q show individual mineralization traces of selected compounds used in the study. For many compounds (e.g. humic acid, phenol, oxalic acid, tannic acid, formic acid, KHP, naphthylamine) as well as for euent and groundwater samples, the production of CO2 peaked within 10 min, although signi®cant production was observed to continue, in most instances, for more than an hour. For other compounds (e.g. glucose and glycerol) an initial delay in excess of 15 min occurred before signi®cant CO2 production was observed. This generation of CO2 then continued for over 1 h. Saturated compounds lacking pi bonds (e.g. adamantane, tetrabutyl ammonium chloride) did not show ``peak'' mineralization, but a rather slow and continuous ramping of the CO2 levels over the course of the experiment. The oxidation of caeine, in particular, displayed an interesting mineralization pro®le (Fig 5b). Caffeine consistently demonstrated an initial CO2 maximum within the ®rst 10 min and a second maximum after 30±90 min, depending on the oxidant and the initial pH. Because it was so distinctive, there was a suspicion that something unusual like a temporary ¯ow interruption or a pressure change in the air supply might have caused these features. Duplication of the plot, however, proved otherwise. The possibility that the bimodal CO2 pro®le may have been caused by pKa changes during the course of oxidation, which is known to occur when carbon or in this case the ring nitrogens of caeine become oxidized, was discounted when the pattern persisted in acidi®ed standards (pH R 3). What was most notable in subsequent runs, however, were the variations in the absolute and relative times of the CO2 maxima when conditions were modi®ed (Fig. 5b). Although intermediates contributing to the observed behavior were not identi®ed, it may be reasoned that the plots re¯ected two distinct phases in caeine oxidation over time. The oxidation of discrete carbons on the caeine ring is one possibility. Bimodal plots may also have resulted from the generation of partially oxidized intermediates which then underwent further, delayed oxidations or decarboxylation to produce a second concentration spike. Similar multi-modal CO2 distribution patterns were also observed in some secondary euent samples. Deciphering the mechanisms which promoted this course of mineralization in caeine or in treatment plant euent samples was beyond the scope of the present study; however, it might be bene®cial to examine the details of this process in the future. Plot attributes may be a useful tool for gauging optimum contact times, or predicting refractory/ non-refractory properties of DOC in chemically uncharacterized waters. It is very likely that the
observed time maxima and pro®le variations of the plots are closely linked to structural and functional features of molecules studied. Aromatic compounds showed an immediate exponential increase in CO2 generation with a correspondingly rapid decrease in concentration; this initial phase was then followed by a slow asymptotic decrease in CO2 levels. In addition, the ozone and ozone:peroxide plots of the pi bonded compounds were, in most cases, virtually superimposable and had insigni®cant dierences in total generated CO2 (Table 1). Conversely, saturated hydrocarbons which are known to be refractory to ozone showed measurably better performance with ozone:peroxide. Their plots also lacked the well de®ned CO2 maxima of pi bonded compounds and had relatively lower overall CO2 concentrations. Another distinctive feature of these plots is the rate at which non-refractory compounds, after the initial concentration maximum, attained a steady state for CO2 production, as was evidenced by the slow, asymptotic decreases in concentration. This relatively low, inconsequential CO2 loss appears to be occurring in spite of persistently high TOC residuals (Fig. 5a). The observed behavior is not likely to be an artifact of changes in the partition coecient as oxidations progressed, because reactions of neutral unbuered samples invariably result in pH decreases that favor CO2 desorption in the latter stages of oxidation. Such responses were only observed for highly refractory compounds (Fig. 5c). There is a possibility, however, that pH reduction as the reactions progressed suppressed the production of hydroxyl radicals. Replication of the general pro®les with acidi®ed samples may suggest that the impact of pH changes at least under conditions employed for this study, may have played only a minor role in the observed mineralization dynamics. Because residual TOCs were approximately seventy percent of the initial values, even for the more reactive compounds, it was remarkable that the latter stages of these oxidations did not always show incremental increases in CO2 concentration as was evident for the refractory compounds (Fig. 5c). One possible explanation for this may be that Eact of Primary oxidation processes is <
FT-IR study of DOC mineralization
rated compounds, undergo a relatively rapid initial oxidation which yields small quantities of CO2 and other transitory products. Following this stage, the oxidized intermediates, and/or their progressively oxidized derivatives, then appear to persist in solution, resist mineralization, and contribute to TOC residuals. Formic, oxalic and acetic acid, known ozonation end-products, like other refractory compounds studied, exhibited higher mineralization with ozone:peroxide. When a sonicated suspension of paraformaldehyde, an oligomer of formaldehyde, was tested the reaction showed an immediate and unexpected high level of CO2 production. This result is surprising since formaldehyde is a known by-product of advanced oxidation processes. Moreover, polymerized formaldehyde would be expected to have relatively lower overall oxidation rates, because polymers typically exhibit diminished reactivity when compared to their monomeric form. The discovery that a partially soluble suspension of a polymerized by-product as well as other ozonation ``end-products'' yielded to further oxidation with both ozone and ozone:peroxide, while signi®cant levels of these and other functionally similar constituents appear to persist in solution and escape oxidation, after prolonged contact with advanced oxidants, may suggest something about the nature of the aqueous micro-environment. The existence of any DOC which might be refractory to complete oxidation with advanced oxidants is remarkable. Mineralization of organic carbon through non-combustive oxidation is well supported (Horie and Moortgat, 1991; Pelizzetti et al., 1989). The reaction of ozone with alkenes, in particular, has been the subject of numerous studies (Nangia and Benson, 1980; Grosjean et al., 1994; Grosjean and Grosjean, 1997). The major features of this reaction's mechanism are now well understood (Harding and Goddard, 1978; Atkinson et al., 1997; Criegee, 1975; Niki et al., 1987). The reaction of O3 with alkenes involves an initial addition to the pi bond of the alkene to form a primary ozonide. This intermediate then rapidly decomposes to form two sets of carbonyls plus biradical products. There is also evidence that hydroxyl radicals are formed, via autocatalytic ozone decomposition, as part of this process (Grosjean and Grosjean, 1997; Horie and Moortgat, 1991). These hydroxyl radicals, having a redox potential of 2.8 V, are suciently reactive to permit the complete oxidation of organic molecules, after relatively short contact times. Radical reaction rates are typically in the range 108±1010 mÿ1 sÿ1 (Buxton et al., 1988), and their reactions with hydrocarbons are four to ®ve orders of magnitude faster than equivalent reactions with molecular ozone (Hoigne and Bader, 1976, 1983a, 1983b). By even the most conservative estimation such reactivity should lead to the non-selective, diusion controlled oxidation of a broad range of organic
4045
compounds, even those which are known to be refractory to oxidants such as chlorine, chloramine and permanganates. These reactions which depend on the creation and utilization of hydroxyl radicals are referred to as ``advanced oxidation'' (Glaze et al., 1982; BeltraÂn et al., 1997). In theory, organic molecules in contact with advanced oxidants should be oxidized to carbon dioxide, water and the inorganic ions of any non-hydrogen substituents. Ozone, alone or in combination with hydrogen peroxide, creates hydroxyl radicals during the course of oxidation (Horie and Moortgat, 1991; Bailey, 1972, 1978) and should therefore be capable of mineralizing dissolved organic carbon. Eect of solute/solvent interactions It was evident from this study, however, that mineralization played an insigni®cant role in DOC removal. Although complete oxidation of organics by the powerful hydroxyl radical should be a facile, diusion controlled process, the eciency of the reaction was surprisingly low, even when measured by TOC elimination. The observed results may be caused by dynamics which prevent oxidations from reaching completion in aqueous matrices. Inherent increases in the polarity of intermediates as oxidation progresses boost hydrophilicity which, in turn, increase solvation. This process may serve to shield intermediates, while concomitant decreases in the electron densities further retard oxidation (Hynes and Baer, 1985; Weaver, 1992). Unlike the gas phase environment where polar reactants can collide and react with relative ease, the aqueous environment allows H2O molecules which possess a high dielectric constant EH2O 81:7Eo to surround and form weak complexes with highly oxidized, polar intermediates. These solvating interactions can be signi®cant depending on the nature of the charges and the inherent ability of the solvent to coordinate with the ions involved (Rips et al., 1988; Wolynes, 1987). The coordinated, ordered hydration spheres which result can present a sizable barrier to oxidants, and substantially increase the Eact of Secondary oxidation processes. The overall solvation energy, excluding entropy and hydrogen bonding eects, could be represented by the sum of the ion-dipole, dipole-dipole, induced dipole-dipole and instantaneous dipole-induced dipole interactions (Equation (1)) (Jae, 1963; Huheey, 1978):
ÿjZ2 jm ÿ2m1 m2 4pr 2 E0 4pr3 E0 1 Z 2a ÿm 2 a ÿ 2 r4 r6
E
1
where E equals the enthalpy of solvation (hydration) derived from the interaction between a charged solute and a polar solvent, m represents the dipole moment, r is the distance between the ion
4046
Steve A. Carr and Rodger B. Baird
and the molecular dipole, a represents inherent polarizability of neutral species, Eo is the dielectric constant of the solvent, and Z equals the ionic charge. This solvation phenomena, in part, might explain why even the addition of excess oxidant was observed to have negligible eects on CO2 yields. The presence of a solvated ``cage'' or an enlarged hydration sphere could sterically shield intermediates and nullify the reactivity of even the most powerful oxidant. Oxidation rates were observed to be measurably inhibited when ``non-refractory'' compounds were enclosed in the apolar polysaccharide cavity of b-cyclodextrin (b-CD). Although these inclusion complexes of b-CD were not perfect models of solvation, the experiments showed unequivocal perturbation in bleaching and mineralization rates in ozone when compared to non-complexed standards. Analogous dynamics may also occur in aqueous matrices as a result of solvation and contribute to the attenuated reactivities observed, especially in the latter stages of disinfection. Based on this postulate, it is possible that polarity and solvation potential may be important determinants of a DOC's refractory nature. It may also explain why ozonation appears to be ineective at removing a large portion of DOC from reclaimed wastewater and surface waters. These waters are likely to contain hydrophillic fractions of highly polar, oxygenated organics derived from natural oxidation processes (Quanrud et al., 1996). In light of this hypothesis, it would be interesting to study photocatalytic mineralization of highly oxygenated substances using conditions equivalent to those described in this paper. Heterogeneous oxidations using TiO2 which have been shown to eectively mineralize DOC (Pelizzetti et al., 1989; Pichat et al., 1990; Pramauro et al., 1997), may eectively circumvent the proposed solvation barriers by adsorbing the organic species directly onto reactive surfaces where the OH oxidants are generated. Similarly, solvation is unlikely to inhibit photolysis when highly oxygenated DOCs are irradiated. Easier access of the radical to partially oxidized intermediates along with possible assistance from photoactivation may result in synergies which facilitate the higher levels of mineralization observed when photocatalytic systems are employed. The results of such studies, if performed, could provide important clues as to why adsorption appears to be so critical to biological and chemical treatment processes, and could further suggest changes which might enhance mineralization, in advanced oxidation processes. CONCLUSION
The importance of oxidative mineralization as a mechanism for TOC removal during disinfection must be determined to fully utilize AOPs in water
and wastewater treatment. Understanding the full impact of applying AOPs is critical; ®rst, to determine how eective a remedy these much vaunted oxidation processes are to solving existing problems, and secondly, to assess how well they might facilitate compliance with future regulatory changes. Though these processes have been shown to surpass the level of disinfection aorded by traditional techniques and improve biodegradability of some fractions of DOC, in this study, it was determined that the use of ozone, or ozone in combination with hydrogen peroxide to chemically eect TOC reduction may have practical limitations, when applied to dilute aqueous solutions containing both refractory and degradable organic carbons. Reactions which were determined to possess very favorable thermodynamics were found, by an indirect measure, to be complicated by solute/solvent interactions which may have served to inhibit oxidations. AcknowledgementsÐThe authors gratefully acknowledge the diligent eorts of following individuals: Andrea Smith, for all the help she provided with Excel, Jorge Garcia for analyzing TOCs, Tom Kelly for his assistance in sampling, Richard Jackson for his helpful programming suggestions, Ann-Marie Caro and Lavern Tamoria for reading this manuscript and ®nally the Sanitation Districts of Los Angeles County for supporting this work.
REFERENCES
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Water Chlorination: Environmental Impact and Health Eects, Vol. 2. Ann Arbor Science Publishers, Inc, Ann Arbor, MI. Kogelschatz U. and Baessler P. (1987) Determination of nitrous oxide and dinitrogen pentoxide concentrations in the output of air-fed ozone generators of high power densities. Ozone: Science and Engineering 9, 287±297. Kuo P. P., Chian E. S. K. and Chang B. J. (1978) Identi®cation of end products resulting from ozonation and chlorination of organic compounds commonly found in water. Environ. Sci. Technol. 11, 1177±1181. Langlais B., Reckhow D. A. and Brink D. R. (eds) (1991) Ozone in Wastewater Treatment Application and Engineering. AWWA Research Foundation, Lewis Publishers, Chelsea, MI. Legube B., Langlais B. and Dore M. (1980) Reactions of ozone with aromatics in dilute aqueous solution: reactivity and biodegradability of oxidation products. Prog. Water Technol. 12, 553±570. Mackay D. and Leinonen P. J. (1975) Rate of evaporation of low solubility contaminants from water bodies to atmosphere. Environ. Sci. Technol. 9, 1178. Miltner R. J., Skukairy H. M. and Summers R. S. (1992) Disinfection by-product formation and control by ozonation and biotreatment. J. of AWWA 84(11), 53±62. Nangia P. S. and Benson S. W. (1980) Thermochemistry and kinetics of ozonation reactions. J. Am. Chem. Soc. 102, 3105. Narkis N. and Schneider-Rotel M. (1980) Evaluation of ozone induced biodegradability of waste-water treatment plant euent. Water Res. 14, 929±939. Niki H., Maker P. D., Savage C. M., Breitenbach L. P. and Hurley M. D. (1987) FTIR spectroscopic study of the mechanism for the gas-phase reaction between ozone and tetramethylene. J. Phys. Chem. 91, 941±946. Pelizzetti E., Minero C., Maurino V., Sclafani A., Hidaka H. and Serpone N. (1989) Photocatalytic degradation of nonylphenol ethoxylated surfactants. Environ. Sci. Technol. 23, 1380±1385. Peng J., Bewtra J. K. and Biswas N. (1994) Volatilization of selected organic compounds from quiescent water. J. Environ. Engng 120(3), 662. Peng J., Bewtra J. K. and Biswas N. (1995) Eect of turbulence on volatilization of selected organic compounds from water. Water Environ. Res. 67(1), 101. Pichat P., D'Oliveira J.-C. and Al-Sayyed G. (1990) Photodegradation of 2- and 3-Chlorophenol in TiO2 aqueous suspensions. Environ. Sci. Technol. 24, 990±996. Pramauro E., Prevot A. B. and Brizzolesi G. (1997) Photocatalytic degradation of carbaryl in aqueous solutions containing TiO2 suspensions. Environ. Sci. Technol. 31, 3126±3131. Rips I., Klafter J. and Jortner J. (1988) J. Chem. Phys. 89, 4288. Roberts P. V., Dandliker P. and Matter-Muller C. (1983) Volatilization of organic pollutants in wastewater treatment-model studies. EPA-R-806631, US EPA, Munic. Environ. Res. Lab. Cincinnati Ohio. Smith J. H., Bomberger D. C. and Haynes D. L. (1980) Prediction of the volatilization rates of high-volatility chemicals from natural water bodies. Environ. Sci. Technol. 14, 1332. van der Kooij D., Hijnen W. A. M. and Kruithof J. C. (1989) The eects of ozonation, biological ®ltration, and distribution on the concentration of easily assimilable organic carbon in drinking water. Ozone: Science and Engineering 11(3), 297±311. Volk C., Roche P., Joret J.-C. and Paillard H. (1997) Comparison of the eect of ozone, ozone±peroxide system and catalytic ozone on biodegradable organic matter of a fulvic acid solution. Water Res. 31(3), 650±656. Weaver M. J. (1992) Dynamical solvent eects on acti-
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Wat. Res. Vol. 34, No. 16, pp. 4036±4048, 2000 7 2000 Elsevier Science Ltd. All rights reserved Printed in Great Britain 0043-1354/00/$ - see front matter
www.elsevier.com/locate/watres
MINERALIZATION AS A MECHANISM FOR TOC REMOVAL: STUDY OF OZONE/OZONE±PEROXIDE OXIDATION USING FT-IR STEVE A. CARR* and RODGER B. BAIRD County Sanitation Districts of Los Angeles County, San Jose Creek Water Quality Laboratory, 1965 S. Workman Mill Road, Whittier, CA 90601-1415, USA (First received 1 July 1999; accepted in revised form 7 January 2000) AbstractÐFT-IR spectroscopy was used to study the course of abiotic mineralization in samples treated with ozone and ozone:peroxide. The importance of oxidative mineralization as a mechanism for dissolved organic carbon (DOC) removal during disinfection must be determined if advanced oxidation processes (AOPs) in water and wastewater treatment is to be fully exploited. Understanding the full impact of these applications is critical for determining how eective a remedy AOPs are for eliminating disinfection by-product formation and for assessing how well these applications might facilitate compliance with future regulatory changes. Although these processes have been shown to surpass the level of disinfection aorded by traditional techniques and improve biodegradability, this study found that the use of ozone, or ozone in combination with hydrogen peroxide, to chemically eect TOC reduction may have practical limitations when applied to dilute aqueous solutions containing both refractory and degradable dissolved organic carbons. Reactions possessing very favorable thermodynamics were found, by an indirect measure, to be complicated by solute/solvent interactions which may have served to inhibit oxidation. It was also shown that ``close to real time'' CO2 plots generated using highly sensitive FT-IR techniques may have some practical utility in assessing mineralization potential and optimizing disinfection processes. 7 2000 Elsevier Science Ltd. All rights reserved Key wordsÐTOC, ozonolysis, mineralization, b-cyclodextrin, infrared, advanced oxidation, hydroxyl radical, disinfection, solvation eect
INTRODUCTION
When conventional water treatment techniques are employed, dissolved organic matter can be removed from water through a combination of biological, physical and chemical means (Weber, 1972; Jolley et al., 1978; Brunet et al., 1982). These processes are known to interact and exhibit complex, changing synergies as treatment progresses. Of the three mechanisms, biodegradation is generally acknowledged to be the primary removal process, for both natural organic matter (NOM) and euent organic matter (EfOM) in treated wastewaters. This is followed by physical adsorption processes which include granular activated carbon, coagulation and settling. Chemical oxidation, by comparison, is believed to make only minor, direct contributions to dissolved organic carbon (DOC) removal, when traditional disinfection techniques are utilized. Ozone, the most powerful of the common disinfectants, has been reported to be very eective at *Author to whom all correspondence should be addressed.
reducing disinfection by-product (DBP) formation (Miltner et al., 1992; Anderson et al., 1990), in addition to enhancing ¯occulation and oxidizing taste and odor compounds in water (Ferguson et al.,1990). In Europe, where it has been used successfully as a disinfectant for several decades, experience has shown that ozonation can alter TOC levels, and produce large concentrations of lowmolecular-weight carbon compounds (Kuo et al., 1978; Legube et al., 1980; Glaze et al., 1987; Hacker et al., 1994). More recently, advanced oxidation processes (AOPs) which utilized ozone as one of the components, have been shown to increase the destruction of complex organic constituents which are refractory to other treatment (van der Kooij et al., 1989; Ferguson et al., 1990; Brunet et al., 1982). Partially oxidized by-products of these AOP reactions have also been shown to exhibit greater biodegradability (Amy et al., 1996; Duguet et al., 1985; Narkis and Schneider-Rotel, 1980; Legube et al., 1980) and have higher adsorption coecients (Janssens et al., 1984), probably due to their increased polarities. Although oxidative mineralization is sus-
4036
FT-IR study of DOC mineralization
4037
Fig. 1. Ozonolysis apparatus.
pected to be a factor in the observed reduction of DBP and TOC in treated waters (Xu et al., 1989; Adams et al., 1997; Volk et al., 1997; Caprio et al., 1987), there have been few investigations to selectively study its mechanisms and its total contribution to TOC removal. This paper will evaluate the ability of ozone and ozone±peroxide to chemically remove dissolved, non-purgeable organics in water through direct oxidative mineralization. MATERIALS AND METHODS
Experiments were carried out at room temperature in a batch mode. Bench scale ozonation was performed using a 2.2-l Pyrex column (1.07 m 54 cm) ®tted with a sintered glass inlet diuser (Fig. 1). The reactor column was constructed of glass; all other ®ttings and tubing were stainless steel or Te¯on. Flow rates to the column were controlled using an in-line pressure regulator in series with adjustable needle valves. The ¯ow rate was measured using a J&W digital ¯ow meter. The ¯ow rate of the ozonized gas from the generator was adjusted to 433 ml/min. This rate permitted a minimum of three volume exchanges in the IR gas cell and one complete displacement of the apparatus dead volume1 per infrared spectra. Infrared spectra were obtained on reaction products using a Mattson Galaxy 5000 FT-IR spectrophotometer equipped with a DTGS detector. A cylindrical glass Flow-Thru Gas Cell (110 30 mm) with Te¯on stopcock, Viton O-ring seals and two CaF2 IR-Transmission Windows (38 6 mm), having a transmission range from 1100 to 50,000 cmÿ1 (9.0±0.20 mm) was used to analyze the gas-
1
(apparatus dead volume = column head volume+tubing volume+gas cell volume).
eous reaction euent. CaF2 windows were selected because of their resistance to water, acids and alkalies. These properties are important because the cell windows are exposed to a continuous stream of moist air which may contain traces of oxidized by-products. Windows were frequently polished between runs, using aluminum oxide powder (<50 mm) and water to remove any surface build-up. Infrared spectra were acquired using the following parameters: starting wavenumber 1350 cmÿ1; ending wavenumber 4000 cmÿ1; resolution of 4.0 cmÿ1; number of scans 64; gain 1; forward velocity 10.0 KHz; reverse velocity 10.0 KHz; time per spectra 1.23 min; apodization usedÐtriangulation. All compounds used in this study were reagent grade or better, and were purchased from Aldrich, Baker or Mallinckrodt. The b-cyclodextrin (b-CD) was donated by Cerestar, Inc. Standards were prepared in a 20-l pyrex carboy at 10 mg/l concentrations, except when solubility was a limiting factor, using Barnstead Nanopure Type II water. After preparation, standards were aerated to aid dissolution, provide mixing and remove residual bicarbonate. In general standards were not acidi®ed or buered prior to oxidation (except where speci®ed). Well waters, groundwater and plant euent samples were acidi®ed to pH R 3 before aeration. Portions of the standard not used on the day of preparation were capped and refrigerated at 48C. Prior to the start of a run, samples were purged completely of carbon dioxide using CO2-free compressed air from a Balston Model 7562 FTIR Purge Gas Generator. The ozonation apparatus was cleaned by ¯ushing with DI water and rinsing, with the standard to be tested, for a minimum of 5 min. After cleaning, the column was ®lled with sample using a FMI pump. A continuous CO2-free air purge was maintained through the apparatus, before and after ®lling, to remove residual CO2 and prevent ambient contamination. For this study, ozone was generated from air or pure oxygen using a PCI Ozone Generator (Model GL-1, West Caldwell New Jersey). The ozone generator was turned on
4038
Steve A. Carr and Rodger B. Baird
for a minimum of 15 min prior to the start of a run to allow the output to stabilize. Ozone concentration in the feedgas [ave. concentration with O2349 mg/l, air 330 mg/ l] and the o-gas streams were measured using a PCI ozone monitor (Model LC-400, West Caldwell, NJ). The ozonized gas stream from the generator was fed continuously through the bottom of the reactor as micro-bubbles for the duration of a run. Sample mixing was accomplished solely via this process. Earlier dye studies determined that the turbulence and agitation resulting from aeration of the reactor column provided adequate mixing. A DI water blank was determined to have a dissolved ozone concentration between 5 and 8 mg/l with an average value of 07 mg/l, using the indigo blue method. The dis-
solved ozone concentration of a plant euent sample was 5 mg/l after 5.0 min contact time. The euent gas stream from the reaction column was conveyed via a 1/40 Te¯on tubing to the gas ¯ow cell in the FT-IR sample chamber. Blank background spectra of the Gas Cell were collected using CO2-free air from the Balston generator. The o gas from the column was then monitored by infrared for CO2 depletion. A column background spectrum was collected when the CO2 infrared absorbance band disappeared. This reference was stored in instrument memory just prior to the addition of the O3 oxidant, and subtracted from all subsequent spectra collected during oxidation. For ozone:peroxide experiments, 3.64 ml of a 3000 mg/l hydrogen peroxide standard was
Fig. 2. (a) Sample IR spectrum from study. (b) Sample IR spectrum with expanded region.
FT-IR study of DOC mineralization added via syringe at the start of the run, or continuously for the duration of a run using an Altec Model 110A HPLC pump connected to the column sample inlet by 1/80 Te¯on tubing (see Fig. 1.). The total concentration of hydrogen peroxide added was 5.0 mg/l. To begin an analysis, the gaseous air purge to the contact column was switched to the ozonated gas stream from the generator using a four-way stainless steel ball valve. The FT-IR was started simultaneously, initiating data collection. The FT-IR spectrophotometer was programmed to continuously collect and store multiple, incremented infrared spectra, for the duration of a run (see Fig. 2.). In addition, the Macro allowed the integrated areas for the CO2 infrared absorption bands to be stored in preassigned output ®les. For this study, 150 infrared spectra were collected over a 3 12 h period to measure CO2 concentrations, and to evaluate other constituents in the ozonated air stream after column contact. Carbon dioxide concentration was calibrated using a Scotty III Certi®ed Standard having a CO2 concentration of 408 mg/l22%, in nitrogen. A detection limit of R0.25 ppb was calculated for CO2 using serial dilutions of this certi®ed standard. The absorbance area between 2410.6 and 2281.4 cmÿ1 was used to calculate CO2 concentration (mg/l). These values were then plotted vs time using Microsoft Excel 5.0. The total evolved CO2 for a given run was determined from the integrated plots. No adjustment was made for evaporative losses which were estimated to be 32% during runs. A NaHCO3 standard spiked into an acidi®ed DI water column blank yielded >78% theoretical CO2 recovery. The b-cyclodextrin inclusion complexes with caeine, methylene blue and methyl orange were prepared in a 40 ml vial by adding 1.0 mmol of the respective reagents to 2.0 mmol slurry of b-cyclodextrin in 20 ml of DI water. The mixtures were stirred overnight prior to use. Bleaching experiments were conducted on highly colored methylene blue and methyl orange standards and their inclusion complexes. The times from the beginning of ozonation to complete color loss for the neat and complexed chromophores were compared. Two hundred and ®fty milliliters of water samples were collected in precleaned amber bottles before and after each analysis. Samples were treated with 0.5 g of sodium arsenite (NaAsO2) to remove oxyresidual, then preserved with 2 ml 1:1 H2SO4 and refrigerated. TOC analyses were conducted per Standard Methods' 19th Edition (APHA, AWWA, WEF, 1995) 5310C Heated Persulfate Oxidation Method. The analyses were run on an OI Analytical Instrument Model 1010 equipped with an autosampler OI Analytical Instrument Model 1051. RESULTS AND DISCUSSION
Bench scale experiments were performed on prepared standards, well waters from a local groundwater recharge basin, and secondary and tertiary euent samples from a wastewater reclamation plant. Conditions used in this study approximated an open, batch treatment process. Experiments were conducted with the goal of optimizing contact times to determine the capacity of ozone and ozone:peroxide to oxidatively remove dissolved organics. Knowing what occurs when DOCs are exposed to a stoichiometric excess of ozone or ozone:peroxide over prolonged periods could better de®ne the limits of advanced oxidation processes, and provide useful insights into how mineralization progresses with time.
4039
The study's design approximated pseudo ®rst order kinetics and optimized oxidation by keeping the ozone concentration in excess for the duration of each test. These close-to-ideal conditions were assisted by a continuous purge which forced the reaction's equilibrium to the right and removed bicarbonate ions, a potential hydroxyl radical scavenger, from the matrix. Compounds were chosen to represent the wide range of structure and functionalities which might reasonably be expected in natural and disinfected wastewater samples. Standards for this study were prepared at 10 mg/l, except when solubility was limiting. Initial concentrations for TOC were in the 1±7 mg/l range. The percent carbon mineralized was based on the molarity of carbon in standards relative to the number of moles of CO2 generated during oxidation. The experimental results for ozone and ozone:peroxide oxidation of 28 model compounds are summarized in Table 1 and Figs 3 and 4. Total CO2 concentration for these experiments was determined from integrated plots. These graphs were also evaluated to determine optimum contact times and to study relative mineralization pro®les. No measurable concentration or discernable pro®le changes were observed for CO2 plots when the feed-gas was switched between air and oxygen, even though oxygen provided >60% increase in ozone generation vs air. This suggests that the applied ozone concentration, for these experiments, was not rate limiting at or above the 30 mg/l air-feed levels. It also suggests that once initial solubility limits were attained, ozone mass transfer rates were determined solely by reaction rates, not feed-gas stoichiometry. FTIR quanti®cation of CO2 demonstrated that mineralization was not the primary or even an important mechanism for TOC removal from standards. Even after unrealistically long (03.5 h) contact with the oxidants, only 3% of the DOC, on average, was mineralized when measured as moles of CO2 generated. As can be seen in Fig. 3, neither oxidant achieved greater than 06% mineralization. The removal of TOC was about an order of magnitude higher with an average of 31%, based on the TOC concentration in standards before the application of the oxidant. Some loss of organic constituents is known to occur when water or wastewater utilities employ aeration treatments (Bell et al., 1988; Mackay and Leinonen, 1975; Smith et al., 1980; Peng et al., 1994, 1995; Roberts et al., 1983, Langlais et al., 1991). Random comparison of standards before and after air purging 030 min without ozone or ozone:peroxide, showed little or no measurable TOC changes. It is therefore unlikely that the TOC decreases observed were due to any signi®cant air stripping of the unoxidized starting materials. There was also no turbidimetric or particulate organic carbon (POC) formation indicated. The TOC data, then, were somewhat surpris-
4040
Steve A. Carr and Rodger B. Baird
ing, absent physical absorptive or biological processes to account for the carbon loss. TOC reduction might also be observed if partially oxidized intermediates from disinfection were resistant to oxidation under the conditions employed for TOC analysis. However, this appears unlikely. When prepared standards of low molecular weight organic acids and other known ozonation end-products were analyzed for TOC, recoveries were close
to quantitative. A more plausible explanation for the non-mineralized TOC losses may be the direct purging of oxidized intermediates from the reaction column. The persistence of multiple IR bands in the 1653±1750 cmÿ1 region (Fig. 2b) suggests the presence of aldehydes and other carbonyl bearing constituents in the post reaction stream. These products could not be resolved, however, because known by-products such as methyl glyoxal and or-
Table 1. Model compounds studied using ozone/ozone:peroxide oxidation Compound/oxidation method Acetic acid/O3 Acetic acid/OPa Adaman./O3 Adaman./OP t-Butanol/O3 t-Butanol/OP Caeine/O3 Caeine/OP t-CDTA/O3 t-CDTA/OP DEGa/O3 DEG/OP 3,5-DHBA/O3 3,5-DHBA/OP EDTA/O3 EDTA/OP Formic A./O3 Formic A./OP Glucose/O3 Glucose/OP Glycerol/O3 Glycerol/OP Humic A./O3 Humic A./OP Inositol/O3 Inositol/OP KHP/O3 KHP/OP 1-NAM.HCl/O3 1-NAM.HCl/OP 1-NED/O3 1-NED/OP Oxalic A./O3 Oxalic A./OP PDT/O3 PDT/OP PEG/O3 PEG/OP PCP/O3 PCP/OP Phenol/O3 Phenol/OP PMA/O3 PMA/OP Quinaldic A/O3 Quinaldic A/OP Resorcinol/O3 Resorcinol/OP Tannic A./O3 Tannic A./OP Tartaric A./O3 Tartaric A./OP TBAC/O3 TBAC/OP TPM/O3 TPM/OP a
Init. [C] mmol.
[CO2] mmol.
Init. TOC (mg/l)
Final TOC (mg/l)
[TOC]/mmol C
D [C] %
D [TOC] %
Qrfx
0.765 0.765 1.453 1.453 1.208 1.208 0.913 0.913 0.505 0.505 0.861 0.861 1.015 1.015 0.589 0.589 0.456 0.456 0.731 0.731 0.816 0.816 0.681 0.681 0.821 0.821 0.876 0.876 1.222 1.222 1.023 1.023 0.510 0.510 0.899 0.899 0.710 0.710 0.207 0.177 1.432 1.432 1.394 1.394 1.270 1.270 1.221 1.221 0.929 0.929 0.592 0.592 1.339 0.840 1.566 1.566
0.005 0.012 0.002 0.004 0.005 0.016 0.030 0.026 0.025 0.027 0.017 0.036 0.035 0.035 0.018 0.019 0.020 0.017 0.021 0.019 0.030 0.029 0.026 0.019 0.031 0.030 0.022 0.022 0.025 0.031 0.022 0.026 0.016 0.017 0.026 0.026 0.008 0.026 0.007 0.008 0.045 0.045 0.031 0.032 0.033 0.033 0.041 0.039 0.024 0.025 0.019 0.019 0.001 0.015 0.001 0.002
4.18 4.18 7.50 7.50 1.70 1.70 3.78 3.78 4.07 4.07 4.34 4.34 5.23 5.23 3.00 3.00 2.80 2.80 3.10 3.10 4.00 4.00 2.20 2.20 4.26 4.26 3.70 3.70 6.50 6.50 4.83 4.83 2.20 2.20 4.30 4.30 4.73 4.73 0.98 0.98 7.43 7.43 6.51 6.51 6.01 6.01 6.30 6.30 4.48 4.48 3.29 3.29 6.69 6.69 6.51 6.51
3.62 2.83 9.70 7.50 2.01 2.01 0.58 0.64 0.22 0.26 1.89 0.00 0.27 0.09 0.20 0.20 0.11 0.12 0.81 0.25 0.18 0.08 0.49 0.65 0.34 0.48 1.00 0.17 2.40 1.50 1.25 1.41 0.17 0.14 0.74 0.81 4.13 1.86 0.59 0.15 0.79 0.49 1.71 1.65 1.27 0.41 3.40 3.00 1.40 1.80 0.76 0.37 6.52 4.05 1.71 0.17
5.47 5.47 5.16 5.16 1.41 1.41 4.14 4.14 8.06 8.06 5.04 5.04 5.15 5.15 5.09 5.09 6.14 6.14 4.24 4.24 4.90 4.90 3.23 3.23 5.19 5.19 4.22 4.22 5.32 5.32 4.72 4.72 4.31 4.31 4.78 4.78 6.66 6.66 4.76 5.55 5.19 5.19 4.67 4.67 4.73 4.73 5.16 5.16 4.82 4.82 5.56 5.56 4.99 7.97 4.16 4.16
0.70 1.56 0.17 0.27 0.40 1.32 3.27 2.84 5.03 5.32 1.99 4.16 3.44 3.46 3.04 3.23 4.30 3.77 2.89 2.60 3.69 3.52 3.76 2.78 3.83 3.62 2.51 2.55 2.03 2.57 2.18 2.50 3.04 3.39 2.88 2.87 1.14 3.63 3.29 4.68 3.15 3.17 2.20 2.27 2.58 2.58 3.33 3.22 2.56 2.69 3.12 3.19 0.07 1.76 0.08 0.10
13.40 32.30 ÿ29.33 0.00 ÿ18.24 ÿ18.24 84.79 83.02 94.59 93.61 56.45 100.00 94.91 98.34 93.40 93.27 96.07 95.71 73.87 91.94 95.43 97.93 77.73 70.45 92.07 88.78 72.97 95.41 63.08 76.92 74.12 70.81 92.27 93.64 82.88 81.26 12.69 60.68 39.84 84.76 89.42 93.38 73.73 74.65 78.87 93.24 46.03 52.38 68.75 59.82 76.93 88.85 2.54 39.46 73.73 97.47
1.02 0.87 1.15 1.14 1.30 0.98 0.87 0.18 0.29 0.24 1.21 1.67 0.25 0.22 0.04 0.12 0.28 0.05 0.63 0.41 0.25 0.17 0.18 0.25 0.43 0.35 0.19 0.15 0.31 0.30 0.20 0.25 0.87 0.08 0.17 0.13 1.18 1.29 0.12 0.05 0.34 0.34 0.10 0.12 0.27 0.23 0.29 0.26 0.27 0.44 0.55 0.35 1.20 0.81 0.37 0.20
Adaman=1-adamantanamine, t-CDTA=t-cyclohexylenediamine-tetraacetic acid, DEG=diethylene glycol, 3,5-DHBA=3,5-dihydroxybenzoic acid, PDT=3-(2-pyridyl)-5,6-diphenyl,-1,2,4 triazine p,p'-disulfonic acid monosodium salt, KHP=potassium hydrogen phthalate, NAM.HCl=1-naphthylamine.HCl, NED=1-naphthylethylenediamine, PEG=polyethylene glycol (MW 8000), PCP=pentachlorophenol, PMA=1-pyrenemethylamine, TBAC=tetrabutylammonium chloride, TPM=triphenylmethanol, OP=ozone:peroxide.
FT-IR study of DOC mineralization
4041
Fig. 3. Percent carbon mineralized for selected compounds during ozone/ozone:peroxide application.
ganic acids all possess fundamental infrared vibrational frequencies in this region (Atkinson et al., 1997). A white ¯aky build-up on the apparatus tubing was determined to be paraformaldehyde when a
sample of the material was heated above 1408C in an IR gas cell and the resulting infrared spectral bands matched those of formaldehyde. This identi®cation was later con®rmed by dissolving a portion
Fig. 4. Percent TOC change for selected compoundsd during ozone/ozone:peroxide application.
4042
Steve A. Carr and Rodger B. Baird
Fig. 5. The ozone/ozone:peroxide mineralization pro®les of selected model compounds.
of the heated vapors in DI water, derivatizing the solution with 2,4-dinitrophenyl hydrazine, then analyzing the hydrazone derivative by HPLC. These analyses showed unequivocally that loss of formaldehyde was occurring during ozonation, and suggest that formaldehyde loss may have been an important mode for organic carbon removal during oxidation. Similarly, the loss of CO as a by-product of DOC oxidation cannot be ruled out in light of results from previous ozonolysis studies (Niki et al.,
1987; Atkinson and Carter, 1984; Atkinson et al., 1997; Horie and Moortgat, 1991). Quantitation of the CO infrared bands at 2025 and 2240 cmÿ1 was subject to interference from an adjoining O3 band at 2097 cmÿ1 and a N2O band at 2236 cmÿ1 in the IR spectra (Kogelschatz and Baessler, 1987) (Fig. 2a). Mineralization pro®les of ozone and ozone:peroxide The data summarized in Table 1 represent the net
FT-IR study of DOC mineralization
results from 3 12 h long experiments. Although this time may be long relative to real-world disinfection processes, the duration of these experiments provided a unique window to study how AOPs progress over time. Use of an in-line FTIR also allowed the production of CO2 to be evaluated in ``close to real time'' for a range of model compounds during ozone and ozone:peroxide applications. The N2O and O3 concentration in the vented gas was monitored simultaneously as illus-
4043
trated in Fig. 2a. As expected, nitrous oxide was absent when oxygen was used as a feed-gas for ozone generation. Only tentative spectral assignments were made for other functional groups such as C=O, C-H, and -COOH, because of the potential for band overlap. (see IR spectra Fig. 2b). In spite of the minor role CO2 elimination had in the reduction of TOC, the variable features of the CO2 evolution plots have provided revealing insights into the complex nature of these oxidation
Fig. 5 (continued).
4044
Steve A. Carr and Rodger B. Baird
processes (Fig. 5a±d). Figure 5a shows the CO2 generation pro®les for selected model compounds, a tertiary wastewater euent, and groundwater samples. Figure 5b±q show individual mineralization traces of selected compounds used in the study. For many compounds (e.g. humic acid, phenol, oxalic acid, tannic acid, formic acid, KHP, naphthylamine) as well as for euent and groundwater samples, the production of CO2 peaked within 10 min, although signi®cant production was observed to continue, in most instances, for more than an hour. For other compounds (e.g. glucose and glycerol) an initial delay in excess of 15 min occurred before signi®cant CO2 production was observed. This generation of CO2 then continued for over 1 h. Saturated compounds lacking pi bonds (e.g. adamantane, tetrabutyl ammonium chloride) did not show ``peak'' mineralization, but a rather slow and continuous ramping of the CO2 levels over the course of the experiment. The oxidation of caeine, in particular, displayed an interesting mineralization pro®le (Fig 5b). Caffeine consistently demonstrated an initial CO2 maximum within the ®rst 10 min and a second maximum after 30±90 min, depending on the oxidant and the initial pH. Because it was so distinctive, there was a suspicion that something unusual like a temporary ¯ow interruption or a pressure change in the air supply might have caused these features. Duplication of the plot, however, proved otherwise. The possibility that the bimodal CO2 pro®le may have been caused by pKa changes during the course of oxidation, which is known to occur when carbon or in this case the ring nitrogens of caeine become oxidized, was discounted when the pattern persisted in acidi®ed standards (pH R 3). What was most notable in subsequent runs, however, were the variations in the absolute and relative times of the CO2 maxima when conditions were modi®ed (Fig. 5b). Although intermediates contributing to the observed behavior were not identi®ed, it may be reasoned that the plots re¯ected two distinct phases in caeine oxidation over time. The oxidation of discrete carbons on the caeine ring is one possibility. Bimodal plots may also have resulted from the generation of partially oxidized intermediates which then underwent further, delayed oxidations or decarboxylation to produce a second concentration spike. Similar multi-modal CO2 distribution patterns were also observed in some secondary euent samples. Deciphering the mechanisms which promoted this course of mineralization in caeine or in treatment plant euent samples was beyond the scope of the present study; however, it might be bene®cial to examine the details of this process in the future. Plot attributes may be a useful tool for gauging optimum contact times, or predicting refractory/ non-refractory properties of DOC in chemically uncharacterized waters. It is very likely that the
observed time maxima and pro®le variations of the plots are closely linked to structural and functional features of molecules studied. Aromatic compounds showed an immediate exponential increase in CO2 generation with a correspondingly rapid decrease in concentration; this initial phase was then followed by a slow asymptotic decrease in CO2 levels. In addition, the ozone and ozone:peroxide plots of the pi bonded compounds were, in most cases, virtually superimposable and had insigni®cant dierences in total generated CO2 (Table 1). Conversely, saturated hydrocarbons which are known to be refractory to ozone showed measurably better performance with ozone:peroxide. Their plots also lacked the well de®ned CO2 maxima of pi bonded compounds and had relatively lower overall CO2 concentrations. Another distinctive feature of these plots is the rate at which non-refractory compounds, after the initial concentration maximum, attained a steady state for CO2 production, as was evidenced by the slow, asymptotic decreases in concentration. This relatively low, inconsequential CO2 loss appears to be occurring in spite of persistently high TOC residuals (Fig. 5a). The observed behavior is not likely to be an artifact of changes in the partition coecient as oxidations progressed, because reactions of neutral unbuered samples invariably result in pH decreases that favor CO2 desorption in the latter stages of oxidation. Such responses were only observed for highly refractory compounds (Fig. 5c). There is a possibility, however, that pH reduction as the reactions progressed suppressed the production of hydroxyl radicals. Replication of the general pro®les with acidi®ed samples may suggest that the impact of pH changes at least under conditions employed for this study, may have played only a minor role in the observed mineralization dynamics. Because residual TOCs were approximately seventy percent of the initial values, even for the more reactive compounds, it was remarkable that the latter stages of these oxidations did not always show incremental increases in CO2 concentration as was evident for the refractory compounds (Fig. 5c). One possible explanation for this may be that Eact of Primary oxidation processes is <
FT-IR study of DOC mineralization
rated compounds, undergo a relatively rapid initial oxidation which yields small quantities of CO2 and other transitory products. Following this stage, the oxidized intermediates, and/or their progressively oxidized derivatives, then appear to persist in solution, resist mineralization, and contribute to TOC residuals. Formic, oxalic and acetic acid, known ozonation end-products, like other refractory compounds studied, exhibited higher mineralization with ozone:peroxide. When a sonicated suspension of paraformaldehyde, an oligomer of formaldehyde, was tested the reaction showed an immediate and unexpected high level of CO2 production. This result is surprising since formaldehyde is a known by-product of advanced oxidation processes. Moreover, polymerized formaldehyde would be expected to have relatively lower overall oxidation rates, because polymers typically exhibit diminished reactivity when compared to their monomeric form. The discovery that a partially soluble suspension of a polymerized by-product as well as other ozonation ``end-products'' yielded to further oxidation with both ozone and ozone:peroxide, while signi®cant levels of these and other functionally similar constituents appear to persist in solution and escape oxidation, after prolonged contact with advanced oxidants, may suggest something about the nature of the aqueous micro-environment. The existence of any DOC which might be refractory to complete oxidation with advanced oxidants is remarkable. Mineralization of organic carbon through non-combustive oxidation is well supported (Horie and Moortgat, 1991; Pelizzetti et al., 1989). The reaction of ozone with alkenes, in particular, has been the subject of numerous studies (Nangia and Benson, 1980; Grosjean et al., 1994; Grosjean and Grosjean, 1997). The major features of this reaction's mechanism are now well understood (Harding and Goddard, 1978; Atkinson et al., 1997; Criegee, 1975; Niki et al., 1987). The reaction of O3 with alkenes involves an initial addition to the pi bond of the alkene to form a primary ozonide. This intermediate then rapidly decomposes to form two sets of carbonyls plus biradical products. There is also evidence that hydroxyl radicals are formed, via autocatalytic ozone decomposition, as part of this process (Grosjean and Grosjean, 1997; Horie and Moortgat, 1991). These hydroxyl radicals, having a redox potential of 2.8 V, are suciently reactive to permit the complete oxidation of organic molecules, after relatively short contact times. Radical reaction rates are typically in the range 108±1010 mÿ1 sÿ1 (Buxton et al., 1988), and their reactions with hydrocarbons are four to ®ve orders of magnitude faster than equivalent reactions with molecular ozone (Hoigne and Bader, 1976, 1983a, 1983b). By even the most conservative estimation such reactivity should lead to the non-selective, diusion controlled oxidation of a broad range of organic
4045
compounds, even those which are known to be refractory to oxidants such as chlorine, chloramine and permanganates. These reactions which depend on the creation and utilization of hydroxyl radicals are referred to as ``advanced oxidation'' (Glaze et al., 1982; BeltraÂn et al., 1997). In theory, organic molecules in contact with advanced oxidants should be oxidized to carbon dioxide, water and the inorganic ions of any non-hydrogen substituents. Ozone, alone or in combination with hydrogen peroxide, creates hydroxyl radicals during the course of oxidation (Horie and Moortgat, 1991; Bailey, 1972, 1978) and should therefore be capable of mineralizing dissolved organic carbon. Eect of solute/solvent interactions It was evident from this study, however, that mineralization played an insigni®cant role in DOC removal. Although complete oxidation of organics by the powerful hydroxyl radical should be a facile, diusion controlled process, the eciency of the reaction was surprisingly low, even when measured by TOC elimination. The observed results may be caused by dynamics which prevent oxidations from reaching completion in aqueous matrices. Inherent increases in the polarity of intermediates as oxidation progresses boost hydrophilicity which, in turn, increase solvation. This process may serve to shield intermediates, while concomitant decreases in the electron densities further retard oxidation (Hynes and Baer, 1985; Weaver, 1992). Unlike the gas phase environment where polar reactants can collide and react with relative ease, the aqueous environment allows H2O molecules which possess a high dielectric constant EH2O 81:7Eo to surround and form weak complexes with highly oxidized, polar intermediates. These solvating interactions can be signi®cant depending on the nature of the charges and the inherent ability of the solvent to coordinate with the ions involved (Rips et al., 1988; Wolynes, 1987). The coordinated, ordered hydration spheres which result can present a sizable barrier to oxidants, and substantially increase the Eact of Secondary oxidation processes. The overall solvation energy, excluding entropy and hydrogen bonding eects, could be represented by the sum of the ion-dipole, dipole-dipole, induced dipole-dipole and instantaneous dipole-induced dipole interactions (Equation (1)) (Jae, 1963; Huheey, 1978):
ÿjZ2 jm ÿ2m1 m2 4pr 2 E0 4pr3 E0 1 Z 2a ÿm 2 a ÿ 2 r4 r6
E
1
where E equals the enthalpy of solvation (hydration) derived from the interaction between a charged solute and a polar solvent, m represents the dipole moment, r is the distance between the ion
4046
Steve A. Carr and Rodger B. Baird
and the molecular dipole, a represents inherent polarizability of neutral species, Eo is the dielectric constant of the solvent, and Z equals the ionic charge. This solvation phenomena, in part, might explain why even the addition of excess oxidant was observed to have negligible eects on CO2 yields. The presence of a solvated ``cage'' or an enlarged hydration sphere could sterically shield intermediates and nullify the reactivity of even the most powerful oxidant. Oxidation rates were observed to be measurably inhibited when ``non-refractory'' compounds were enclosed in the apolar polysaccharide cavity of b-cyclodextrin (b-CD). Although these inclusion complexes of b-CD were not perfect models of solvation, the experiments showed unequivocal perturbation in bleaching and mineralization rates in ozone when compared to non-complexed standards. Analogous dynamics may also occur in aqueous matrices as a result of solvation and contribute to the attenuated reactivities observed, especially in the latter stages of disinfection. Based on this postulate, it is possible that polarity and solvation potential may be important determinants of a DOC's refractory nature. It may also explain why ozonation appears to be ineective at removing a large portion of DOC from reclaimed wastewater and surface waters. These waters are likely to contain hydrophillic fractions of highly polar, oxygenated organics derived from natural oxidation processes (Quanrud et al., 1996). In light of this hypothesis, it would be interesting to study photocatalytic mineralization of highly oxygenated substances using conditions equivalent to those described in this paper. Heterogeneous oxidations using TiO2 which have been shown to eectively mineralize DOC (Pelizzetti et al., 1989; Pichat et al., 1990; Pramauro et al., 1997), may eectively circumvent the proposed solvation barriers by adsorbing the organic species directly onto reactive surfaces where the OH oxidants are generated. Similarly, solvation is unlikely to inhibit photolysis when highly oxygenated DOCs are irradiated. Easier access of the radical to partially oxidized intermediates along with possible assistance from photoactivation may result in synergies which facilitate the higher levels of mineralization observed when photocatalytic systems are employed. The results of such studies, if performed, could provide important clues as to why adsorption appears to be so critical to biological and chemical treatment processes, and could further suggest changes which might enhance mineralization, in advanced oxidation processes. CONCLUSION
The importance of oxidative mineralization as a mechanism for TOC removal during disinfection must be determined to fully utilize AOPs in water
and wastewater treatment. Understanding the full impact of applying AOPs is critical; ®rst, to determine how eective a remedy these much vaunted oxidation processes are to solving existing problems, and secondly, to assess how well they might facilitate compliance with future regulatory changes. Though these processes have been shown to surpass the level of disinfection aorded by traditional techniques and improve biodegradability of some fractions of DOC, in this study, it was determined that the use of ozone, or ozone in combination with hydrogen peroxide to chemically eect TOC reduction may have practical limitations, when applied to dilute aqueous solutions containing both refractory and degradable organic carbons. Reactions which were determined to possess very favorable thermodynamics were found, by an indirect measure, to be complicated by solute/solvent interactions which may have served to inhibit oxidations. AcknowledgementsÐThe authors gratefully acknowledge the diligent eorts of following individuals: Andrea Smith, for all the help she provided with Excel, Jorge Garcia for analyzing TOCs, Tom Kelly for his assistance in sampling, Richard Jackson for his helpful programming suggestions, Ann-Marie Caro and Lavern Tamoria for reading this manuscript and ®nally the Sanitation Districts of Los Angeles County for supporting this work.
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