Advanced oxidation of phenolic compounds

Advances in Environmental Research 4 Ž2000. 233᎐244

Advanced oxidation of phenolic compounds R. Alnaizy, A. AkgermanU Department of Chemical Engineering, Texas A & M Uni¨ ersity, College Station, TX 77843, USA Accepted 18 May 2000

Abstract Phenol degradation with a UVrH2 O2 advanced oxidation process ŽAOP. was studied in a completely mixed, batch photolytic reactor. The UV irradiation source was a low-pressure mercury vapor lamp that was axially centered and was immersed in the phenol solution. The effects of hydrogen peroxide dosage, initial phenol concentration, H2 O2rphenol molar ratio, pH, and temperature have been investigated. The experimental results indicate that there is an optimum H2 O2rphenol molar ratio in the range of 100᎐250. A sufficient amount of hydrogen peroxide was necessary, but a very high H2 O2 concentration inhibited the photoxidation rate. The second-order reaction rate constants were inversely affected by the initial phenol concentration. No pH effect was observed in the pH range of 4᎐10. A detailed reaction mechanism was proposed. The reaction products include hydroquinones, benzoquinones, and aliphatic carboxylic acids with up to six carbon atoms. A kinetic model, which employs the pseudo-steady state assumption to estimate hydroxyl radical concentration and assumes constant pH was developed to predict phenol oxidation kinetics and product distribution. 䊚 2000 Elsevier Science Ltd. All rights reserved. Keywords: Ultraviolet radiation; Phenol; Photolysis; Hydrogen peroxide; Photoxidation

1. Introduction Chemical oxidation processes are widely used to treat drinking water, wastewater, and groundwater contaminated with organic compounds. Direct oxidation of aqueous solutions containing organic contaminants can be performed under a variety of conditions ranging from ambient conditions to supercritical water oxidation at very high temperatures and pressures. Oxidation at mild conditions by reactive species, such as hydroxyl radicals generated by UV radiation in the

U

Corresponding author. Tel.: q1-978-845-3375; fax: q1978-845-6446. E-mail address: [email protected] ŽA. Akgerman..

presence of an oxidant, such as ozone or hydrogen peroxide, is referred to as an advanced oxidation process ŽAOP.. AOPs are attractive alternative technologies for destroying toxic organic contaminants. AOPs are studied in different combinations; ozone with ultraviolet radiation; ozone with hydrogen peroxide; ozonerhydrogen peroxide with ultraviolet radiation; hydrogen peroxide with ultraviolet radiation; and ozone at high pH. Ultraviolet photolysis combined with hydrogen peroxide ŽUVrH2 O2 . is one of the most appropriate AOP technologies for removing toxic organics from water because it may occur in nature itself. This process involves the production of reactive hydroxyl radicals ŽOH . . that are ultimately capable of mineralizing organic contaminants. This oxidation may occur via one of three general pathways: Ž1. hydrogen

1093-0191r00r$ - see front matter 䊚 2000 Elsevier Science Ltd. All rights reserved. PII: S 1 0 9 3 - 0 1 9 1 Ž 0 0 . 0 0 0 2 4 - 1

234

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

abstraction; Ž2. electron transfer; and Ž3. radical addition ŽMasten and Davies, 1994.. Phenols are one of the most abundant pollutants in industrial wastewater, i.e. chemical, petrochemical, paint, textile, pesticide plants, etc. They serve as intermediates in the industrial synthesis of products as diverse as adhesives and antiseptics. Chlorinated phenols, which form during chlorination of water, are reported to resist biodegradation ŽLa Grega et al., 1994.. The objective of this study was to investigate the feasibility of treating phenol contaminated water by the UVrH2 O2 process. We performed direct ultraviolet photolysis, hydrogen peroxide oxidation, and ultraviolet photolysis combined with hydrogen peroxide oxidation of model phenol solutions. We performed experiments at various initial phenol and H2 O2 concentrations to investigate the effects of their initial concentration on the oxidation rates. From the observed phenol, intermediates, and final product concentration time profiles, we calculated the rate of each reaction in the proposed mechanism. Finally, we proposed a kinetic model and determined the reaction rate parameters for phenol decomposition.

2. Methodology 2.1. Reactor configuration We performed batch experiments in a 6.5-cm diameter cylindrical Pyrex glass jacketed reactor with a total volume of 310 ml. Fig. 1 is a schematic representation of the apparatus used in this study. At the top, the reactor had inlets for feeding reactants, and ports for measuring temperature and withdrawing samples. The reactor was open to air with a magnetic stirrer placed in the bottom to provide proper mixing. The irradiation in the photoreactor was obtained by a low-pressure mercury lamp ŽPCQ lamp, UVP, Inc.. positioned and immersed in the solution in the center of the reactor. The lighted length of the lamp was 63.5 mm with a quartz sleeve diameter of 9.5 mm. A 24᎐40 ground-glass taper joint supported the lamp. It emits approximately 90% of its radiation at 254 nm with a 15-W power input. Typical intensity of the lamp at 4.84 cm2 was 5400 ␮W cmy2 and its body temperature was approximately 60⬚C ŽTable 1.. Because the light source produces heat, in order to conduct experiments at room or controlled temperatures, the reactor was surrounded with a cooling jacket to maintain a constant temperature. Temperature readings were taken using an Omega digital thermometer with an extended probe that was always immersed in the solution.

Table 1 Reactor and light source specifications Reactor specifications Volume Diameter Material Reflection factor

310 65 Pyrex glass 1

ml mm

Light source Lighted length Quartz diameter Radiation Ž) 90%. Power Body temperature Intensity at 4.84 cm2 UV radiation addition rate

63.5 9.5 254 15 60 5400 1.516= 10y6

mm mm nm W ⬚C ␮Wrcm2 Erl ⭈ s

2.2. Materials Phenol, its expected degradation intermediates, and final products were purchased from Sigma Chemicals unless otherwise stated. The phenol stock solution was prepared by adding an appropriate amount of a reagent-grade 90-wt.% solution to deionized-distilled water to obtain a solution of desired concentration. Hydrogen peroxide stock solution was a reagent-grade 35-wt.% solution used as received. Intermediate ring compounds such as 1,2-benzenediol Žcatechol., 1,3-benzenediol Žresorcinol., 1,4-benzenediol Žhydroquinone.,

Fig. 1. Schematic drawing of the apparatus used in this study. 1, reactor; 2, Hg lamp; 3, cooling jacket; 4, temperature measuring inlet; 5, sampling point; 6, magnetic stirrer.

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

and p-benzoquinone were G 98% pure commercial compounds used as received. O-Benzoquinone was unavailable commercially due to its instability. The acids ᎏ muconic, maleic, fumaric, malonic, succinic, oxalic, acetic, and formic ᎏ were ) 98% pure commercial compounds used as received. The selected intermediate and final organic-acid solutions were prepared by weighing the proper amount to give the desired concentration in deionized distilled water. We kept these aqueous solutions refrigerated Žf 5⬚C. and none of the substances, except p-benzoquinone, was hydrolyzed significantly. Solvents Ži.e. methanol and phosphoric acid. were HPLC grade. The pH of aqueous solutions was adjusted using either sodium hydroxide or hydrochloric acid where needed. Distilled and deionized water, which was further purified with a Barnstead Ultrapure mixed-bed cartridge, was used in cleaning and experimentation. Compressed helium and nitrogen gas with better than 99.9% purity was obtained from Bob Smith Gas Products Company, Bryan, Texas, USA.

2.3. Experimental procedures For a standard reaction run, 250 ml of aqueous solution was used. It was prepared by adding a predetermined amount of hydrogen peroxide, as a concentrated solution Ž35 wt.%., to the contaminated water. The initial phenol concentration was in the range of 40᎐500 ppm Ž4.3= 10y4 ᎐5.3= 10y3 M.. High concentrations were good for easy detection, but they can absorb too much of the UV radiation. The molar ratio of H2 O2 to phenol was varied in the range of 0᎐500, which resulted in H2 O2 concentrations up to 0.5 M. The solution was stirred by a magnetic stirrer at approximately 300 rev. miny1 for 15 min to ensure sufficient mixing. It was verified experimentally that hydrogen peroxide does not self-decay during this period. Turning on the ultraviolet light while continuing mixing under ambient temperature and pressure started the reaction, then, a stopwatch was immediately started. The reaction was carried out at neutral pH Ži.e. 6᎐7. except for a few runs to test for pH effects on the reaction rates. Periodically, samples were withdrawn Ž2.5 ml. through the sampling ports using a long hypodermic stainless steel needle then stored immediately in 20-ml screw-cap glass vials for analyses. Before sampling, pH measurement was taken using a Cole Parmer microcomputer, pH-vision model 05669-20 pH-meter, and electrode.

2.4. Analytical procedures Total organic carbon ŽTOC. analyses were performed using an OI Analytical Corporation Model 700 TOC analyzer. Reaction intermediates were identified

235

by a Dionex high-pressure liquid chromatograph ŽHPLC. using two different columns ŽSupelco.. The first column was a general classical reversed-phase, LC-8; 25 cm by 4.6-mm i.d. filled with 5-␮m o.d. particles on a bonded support Žsilica. and had a pore size of ˚ The UV detector was set at 280 nm. The best 120 A. separation was achieved with a programmed solvent gradient as a mobile phase where the eluent was made less polar as time progressed. It started with a methanolrwater ratio of 1:99, with a flow rate of 1.0 ml miny1, then to 100% methanol in 20 min, then to the initial ratio in 5 min. The second column was a SUPELCOGEL-H column for organic acids analysis. The column was 25 cm = 4.6 mm i.d., packed with sulfonated polystyrene divinylbenzene spherical particles of 9-␮m diameter. The UV detector was set at 210 nm. Separation was based on the ion exclusion method, in which the analytes were selectively partitioned between the resin phase and the external aqueous phase. Analyses were performed at low pH with a simple isocratic mode, where the mobile phase composition was maintained constant, 1% phosphoric acid ŽH3 PO4 . with a flow rate of 0.4 ml miny1. All columns were operated at room temperature and the injection volume was 100 ␮l.

2.5. Reaction kinetics In developing the reaction kinetics we assumed: 1. no reaction between phenol and its byproducts to form higher molecular weight intermediates Žwhich was verified by HPLC.; 2. radiation absorption by any substance other than hydrogen peroxide was negligible; and their oxidation rates by UV radiation alone was negligible, hence oxidation by UV alone was negligible. ŽAt 254 nm the extinction coefficient and quantum yield for phenol are 516 My1 ⭈ cmy1 and 0.05 molecule ⭈ photony1, respectively ŽDulin et al., 1986.. We neglected direct photolysis of phenol because of its low quantum yield to hydrogen peroxide and due to experimental verification indicating negligible photooxidation.; 3. organic and inorganic solutes or impurities did not affect the hydrogen peroxide photolysis rate ŽDe Laat et al., 1994.. Hydrogen peroxide was always present in excess Ž) 40:1., therefore, the hydrogen peroxide concentration was assumed constant. The direct photolysis rate of hydrogen peroxide is represented by ru¨ s ⌽ H 2 O 2 f H 2 O 2 I⬚ Ž 1 y eyA t .

Ž1.

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

236

where ⌽H 2 O 2 is the hydrogen peroxide quantum yield, Io is the UV radiation intensity, and f H 2 O 2 is the fraction of radiation that is absorbed by H2 O2 which, based on our second assumption, equals unity. The solution total absorbance Ž At . at the UV radiation source wavelength Ž ␭ . is given by A t ( 2.303 l ␧ H 2 O 2 C H 2 O 2

Ž2.

where A is the absorbance, ␧H 2 O 2 and C H 2 O 2 are the molar extinction coefficient and H 2 O2 concentration, respectively, and l is the effective reactor light path Žannular .. Table 2 shows the reported rate constants for the most widely accepted reaction scheme of hydrogen peroxide photolysis. The concentrations of both radicals ŽOH . and . OOH. were obtained using the pseudo-steady-state hypothesis, which yields w OH . x ss s

2 ⌽ H 2 O 2 I0 Ž 1 y eyA t . k 2 C H 2 O 2 q k 5 C . OOH q

Ž3.

n

Ý k i Ci is1

and HO 2.

ss

s

k 2 COH⭈C H 2 O 2 k 4 COOH q k 5 COH⭈

Ž4.

where ki are reaction rate constants where hydroxyl radicals are consumed in phenol and reaction intermediates and Ci are phenol and reaction intermediate concentrations. In Eq. Ž3., the numerator represents the hydrogen peroxide direct photolysis rate, that is the OH ⭈ production rate, and the denominator represents the hydroxyl radicals’ consumption rate by all the identified species. The reaction rate between species i in the postulated mechanism and hydroxyl radicals is represented by dCi s Ý k j C j COH ⭈y Ý k i Ci COH ⭈ dt j

Ž5.

i

where Table 2 Rate constants for OH Rxn no. 1 2 3 4 5

. formation by H O 2

2

.

H 2 O 2 q h¨ ª 2OH OH q H 2 O 2 ª OOH q H 2 O OOH q H 2 O 2 ª OH q H 2 O q O 2 2 OOH ª H 2 O 2 q O 2 OOH q OH ª H 2 O q O 2

. . .

.

.

.

䢇

䢇

i,j: phenol, identified intermediates and final products; kj : reaction rate constant where substance i is formed; and ki : reaction rate constant where substance i is consumed.

A kinetic model was developed that describes the kinetics of oxidizing phenol-contaminated water by combined ultraviolet radiation and hydrogen peroxide. The steady-state concentration of the OH . radical and parameter estimation of the above equations were obtained using the software packages Excel and SimusolveTM.

3. Results and discussion 3.1. pH effects All phenol advanced oxidation experiments were performed in distilled water solutions. The pH of the solution decreased from 6.5" 0.5 to an average of 3.2" 0.2 due to formation of acidic intermediates. Other runs were made at higher pH, but the oxidation rates were independent of pH. De Laat et al. Ž1994. showed that the efficiencies of UVrH2 O2 processes were not affected by pH below 8, although a decrease was observed for higher pH. Lipczynska-Kochany Ž1993. studied phenol oxidation by the UVrH2 O2 process and observed no significant effects in the pH range from 7.0 to 9.0. On the other hand, phenol degradation and catechol formation decreased rapidly when pH- 7 and pH) 9. This observed decrease was probably due to the fast decomposition of hydroxyl radicals and hydrogen peroxide at high pH ŽChristensen et al., 1982.. However, direct photolysis of phenol was accelerated in alkalized solutions due to the increase in ‘phenolate anions light absorbency’. Beltran et al. Ž1996., Stefan et al. Ž1996. observed that direct photolysis contributions decreased when pH increased from 2 to 7 and then increased to 60% at pH 12. In addition, they reported that the initial rate of acetone removal was independent of pH in the range of 2᎐7. At pH 10, the initial

photolysis

Reaction

.

䢇

Reaction rate constant ŽMy1 ⭈ sy1 .

Source

ruv s 1.52= 10y6 M ⭈ sy1 2.7= 10 7 Ž0.5" 0.09. 8.3= 10 5 8.0= 10 9

This work Buxton et al. Ž1988. Weinstein and Bielski Ž1979. Buxton et al. Ž1988. Elliot and Buxton Ž1992.

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Fig. 2. Phenol oxidation at various H2 O2rphenol ratios and constant initial phenol concentration. Experimental conditions: T s 27⬚C, C0Ph s 2.23" 0.3= 10y3 M.

rate was inhibited. This was explained in terms of hydrogen peroxide dissociation in alkaline media. Also, the fast reaction of hydroxyl radicals with hydrogen peroxide was responsible for the observed decrease in phenol destruction in acidic media. Consequently, photolysis combined with hydrogen peroxide was most effective in neutral solutions. Based on these findings and because our runs were similarly independent of pH Žwithin the range of 6᎐10. all subsequent runs were performed in neutral media ŽpH 6᎐7..

3.2. Phenol oxidation by ultra¨ iolet radiation A few runs were performed using ultraviolet radiation alone to oxidize phenol in aqueous media in the same reactor. The results of one experiment are shown in Fig. 2 at an initial phenol concentration of 210 " 10 ppm Ž2.23= 10y3 M. at 45⬚C. Almost 25% of phenol was oxidized in 1.5 h and more than 45% degradation occurred over a 4-h period due to direct UV photolysis, or hydroxyl radical attack. However, we have also observed that oxidation by UV alone does lead to formation of two or more ring compounds ŽAlnaizy, 1999..

3.3. H2 O2 r phenol molar ratio effects Fig. 2 shows a few runs that were conducted at a constant initial phenol concentration of 210 " 10 ppm Ž2.23= 10y3 M.. Hydrogen peroxide doses were varied between 5.0= 10y3 and 0.5 M to determine the effects of the H2 O2rphenol molar ratio. We have evaluated a very wide range of H2 O2rphenol molar ratios Ž2᎐500.. Hydrogen peroxide enhanced the photolytic degradation of phenol greatly relative to direct photolysis. However, H2 O2 concentrations had two opposing effects on the reaction rate. Increasing the initial hydro-

237

gen peroxide concentration enhanced the oxidation process up to a certain point at which hydrogen peroxide started to inhibit the phenol photolytic degradation. At higher hydrogen peroxide concentrations, Reaction Ž2. in the H2 O2 photolysis mechanism ŽTable 2. became very important and hydrogen peroxide acted as a free-radical scavenger itself, thereby decreasing the hydroxyl radicals concentration. At our operating wavelength and light intensity, we deduced that the optimum H2 O2rphenol molar ratio was between 100 and 250. At a H2 O2rphenol molar ratio of 125, more than 95% of phenol was oxidized in 40 min. At higher H2 O2rphenol ratios Ž) 300., hydrogen peroxide negatively affected the ultraviolet photolytic process despite the higher free-radical production rate. In agreement with Glaze et al. Ž1995., Beltran et al. Ž1996. a significant reduction in the degradation rate was observed, which revealed the inhibitory effects of excess H2 O2 .

3.4. Initial phenol concentration effects Fig. 3 shows results of other runs conducted at a H2 O2rphenol molar ratio of f 200, at 27 " 2⬚C and various initial phenol concentrations. At a concentration of 40 ppm Ž4.26= 10y4 M., approximately 90% oxidation was achieved in 20 min and complete oxidation was achieved in less than 30 min. The photolysis rate decreased significantly when the initial phenol concentration was increased to approximately 500 ppm Ž5.15= 10y3 M.; only 14% was oxidized over a 30-min period. Most previous authors did not investigate the effects of initial concentration on reaction rates. Nearly all investigators have reported a pseudo-first-order reaction rate for UVrH2 O2 oxidation ŽHong et al., 1966; Glaze et al., 1995; Scheck and Frimmel, 1995; Stefan et al., 1996.. Wei and Wan Ž1991. studied the photocatalytic oxidation of phenol in oxygenated solution with suspensions of titanium dioxide powder. They observed that high initial phenol concentrations had a negative effect on the pseudo-first-order reaction rate constant. In other words, phenol photodecomposition decreased with increasing initial phenol concentration. They also observed that the reduction rate of chemical oxygen demand ŽCOD. was much slower than the phenol decomposition rate. In addition, Augugliaro et al. Ž1988. have reported a negative first-order reaction kinetic for phenol photocatalytic decomposition and the initial phenol concentration had a remarkable inhibitory effect on the apparent rate constant. Although the pseudo-first-order reaction kinetics may give a good approximation over a wide range of concentrations, it does not represent the true rate, because the rate constants depend on the initial concentration. Hence, it was necessary to develop a model that would take

238

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Fig. 3. Phenol oxidation at constant H2 O2rphenol ratio f 200, T s 27⬚C and various initial phenol concentration, C0Ph s 0.43᎐5.2= 10y3 M.

into consideration this initial concentration dependency.

3.5. Temperature effects To determine the effect of temperature on reaction rates, several runs were performed at temperatures between 27 and 45⬚C. Fig. 4 shows results of two runs conducted at a H2 O2rphenol molar ratio of 200 and initial phenol concentration of f 480 ppm Ž5.2= 10y3 M.. As expected, heating enhanced the oxidation rate significantly: at the higher temperature, more than 90% oxidation occurred vs. less than 30% oxidation for the other, over the same 90-min period.

4. Reaction mechanism We have identified the following reaction intermediates by HPLC analysis: 1,2-dihydroxybenzene Žcatechol., 1,3-dihydroxybenzene Žresorcinol., 1,4-dihy-

Fig. 4. Phenol photolytic oxidation at constant H2 O2rphenol 0 ratio f 200, initial phenol concentration, CPh f 5.2= 10y3 M, and different temperatures.

droxybenzene Žhydroquinone., 2,5-cyclohexadiene-1,4dione, muconic acid, maleic acid, fumaric acid, oxalic acid and formic acid. The difference between initial carbon concentration Ž t s 0. and the carbon concentration at time Ž t . was assumed to be the CO2 concentration. This was determined by means of TOC analysis. Most of the products listed above are identical to those reported by Devlin and Harris Ž1984., Scheck and Frimmel Ž1995.. The only carboxylic acid that was detected over the entire reaction time was oxalic acid. However, Devlin and Harris Ž1984. reported that glyoxal and glyoxylic acids have the same retention times as oxalic acid and they cannot be separated by HPLC, hence they may also have been present in the oxalic acid amounts we determined. Ogata et al. Ž1983. stated that excited molecules of organics often respond differently than those in ground states and hence photolysis is expected to proceed via a different mechanism than wet oxidation. Nevertheless, phenol photolysis in the presence of hydrogen peroxide followed almost the same path as the Devlin and Harris Ž1984. mechanism for aqueous phenol oxidation with dissolved oxygen.

4.1. Formation of aromatic intermediates The breakdown of hydrogen peroxide to form hydroxyl radicals activates the phenol aromatic ring through a strong resonance electron-donating effect. It is well known that this effect is felt most strongly at the ortho and para positions. This fact was confirmed by our results; hydroquinone and catechol appeared in amounts sufficient to quantify, but a very low yield of resorcinol was detected. Resorcinol was present only in trace amounts, F 2%, and the catecholrhydroquinone ratio was always greater than unity. Our results agree with those of Lipczynska-Kochany Ž1993. who studied flash photolysis of phenol in the presence of hydrogen peroxide and concluded that ortho-hydroxylation Ži.e. catechol formation. was predominant. In our study, these dihydroxybenzenes were produced and observed in the same product distribution under all experimental conditions. Lipczynska-Kochany Ž1993., however, reported that hydroquinone was the main primary product of phenol direct photolysis in dilute degassed aqueous solutions. He also detected 1,2,4- and 1,2,3-trihydroxybenzene Žpyrogallol., though this further hydroxylation reaction on the aromatic ring was of minor yield and was not investigated. If this configuration were maintained in this reaction, that is preferential orthohydroxylation, o-benzoquinone formation should transcend p-benzoquinone concentrations. However, there was no evidence to support this, from this work or others, because: Ž1. benzoquinones, especially ortho, are fairly unstable and are easily cleaved to form aliphatic compounds; Ž2. benzoquinones and dihydroxybenzenes are in redox equilibrium, which may cause

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Fig. 5. Concentration histories of the aromatic intermediates in phenol oxidation by UVrH2 O2 . Experimental conditions: C 0P h s 5.17 = 10 y 3 M; T s 45⬚C; C H 2 O 2 s 0.245 M; H2 O2rphenol molar ratio s 47.

shifting during sampling and HPLC analysis ŽScheck and Frimmel 1995.; Ž3. as mentioned earlier, o-benzoquinone standards were not available commercially, which prevented us from either confirming its presence or following its concentration vs. time. Typical concentration histories of these aromatic substances are shown in Fig. 5. From Fig. 5, it is clear that the ortho-hydroxylation was preferred over the meta or para-substitution. The catechol concentration was approximately 25 times higher than the hydroquinone concentration at the beginning of the reaction and resorcinol was present in a minute amount. In all runs that utilized UV irradiation combined with hydrogen peroxide, all three dihydric phenols were produced. We also detected p-benzoquinone in all runs except in those that were performed at low initial phenol concentration. Neither p-benzoquinone nor resorcinol was observed at low initial phenol concentration, because of the faster reaction rate at low initial phenol concentrations ŽFig. 3.. At low temperatures, keeping all other parameters constant, resorcinol was detected but not p-benzoquinone. This observation may indicate that p-benzoquinone decomposition was very rapid when sufficient hydroxyl radicals were present, even at low temperatures.

239

positively identified. Under all of the experimental conditions used in this study, we detected measurable quantities of the following acids: z, z-muconic, maleic and its isomer fumaric, oxalic, and formic acids. The exceptions were the runs conducted at low H2 O2rphenol molar ratio and at high initial phenol concentration. At high initial phenol concentration, formic acid was observed in a very small concentration in discrete samples, hence, no conclusion could be made about its presence. There were five other unknown peaks. They must have been due to carboxylic andror aliphatic acids because they always eluted within the first few minutes of HPLC analysis. In addition, we know that they were not succinic acid, malonic acid, or acetic acid because none of the unidentified peaks matched these acids when previously used to calibrate standards retention time. Furthermore, they could not be glyoxal and glyoxylic acids because their responses were reported by Devlin and Harris Ž1984. to be similar to that of oxalic acid. Scheck and Frimmel Ž1995. reported the presence of three muconic isomers Ž Z, Z; E, E; and Z, E configurations.; we have confirmed the formation of the first isomer Ž Z, Z . but the others were unavailable commercially to prove their presence. The presence of acrylic acids was unlikely because 3-carbon compound formation was not bound to occur. Based on the acid intermediates and products detected, hydroxyl radical attacks on the double bonds of the unsaturated muconic acid formed maleic and fumaric acid Ž4-carbon acid. as well as oxalic acid Ž2-carbon acid.. Another route is the cleavage of the 1,4-dihydroxybenzene double bond after being hydroxylated to form maleic and oxalic acid ŽScheck and Frimmel, 1995.. However, Devlin and Harris Ž1984. confirmed that acrylic acids were produced but in very low concentrations under all conditions Že.g. excess oxygen or excess phenol.. In any case, these unidentified peaks were readily photoxidized and they disappeared as the reaction progressed. Consequently, their quantification was not possible and

4.2. Formation of carboxylic acids Organic acids were formed and detected as intermediates andror final products under all conditions. The formation of these acids was consistent with pH changes in all experiments. The aqueous solution initial pH was ; 6.8 and within the first 30 min of irradiation, the pH ranged between 4.3 and 4.7. Then slowly over the reaction period, it decreased to approximately 3.2. By means of HPLC analysis using the previously described organic acids column, only five acidic compounds were

Fig. 6. Concentration histories of the carboxylic acids, C0Ph s 1.74= 10y3 M; T s 45⬚C; CH2 O2 s 0.147 M; H2 O2 rphenol molar ratio s 84.

240

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

they were lumped into one term that we called ‘acids’ in the kinetic study. Fig. 6 shows acid concentration histories in UVrH2 O2 oxidation runs at 45 " 2⬚C. Muconic, maleic, fumaric and oxalic acids appeared within the first 10 min of reaction time whereas formic acid appeared after approximately 30 min of reaction time had elapsed. When the initial phenol concentration was doubled, formic acid was not observed and it took maleic acid almost 4 h to start disappearing, though very slowly. This is another clear indication that the initial phenol concentration inhibits the oxidation rate, which we will discuss in detail in Section 5. Oxalic acid was the main reaction product over the entire reaction time in all runs under all conditions. On the basis of this observation and in order to calculate the molar concentration of the lumped-term acids in the kinetics study, we used the molecular weight of oxalic acid as the average molecular weight of the acids to estimate its concentration. The steady rise and the relatively higher concentrations of maleic and oxalic acid over the first 4 h may confirm the hypothesis of Scheck and Frimmel Ž1995., the cleavage of the 1,4-dihydroxybenzene double bond after hydroxylation. They postulated that one double bond of benzoquinone may be hydroxylated then cleaved by the oxidation process yielding maleic acid Žor its isomer fumaric acid. and oxalic acids without going through an intermediate such as muconic acid. This observation would apply more to p-benzoquinone, because the fumaric acid concentration was always much less than that of its isomer maleic acid.

4.3. Total organic carbon (TOC) history Carbon dioxide concentrations were determined by material balances assuming that any unaccounted for carbon mineralizes to carbon dioxide ŽCO2 .. The rate of carbon dioxide formation was calculated by averaging both readings from HPLC analysis and the TOC analyzer, then the difference between TOC and the carbon amount that was initially present in phenol resulted in carbon dioxide production. Also, the carbon in the unidentified intermediates Žacids. was found from the difference between the TOC readings and the carbon that was present in phenol and known intermediate ring compounds. Fig. 7 shows TOC concentration vs. time for several runs at different initial phenol concentrations and temperature. In a typical run, approximately 10 wt.% of the initial carbon present in phenols was converted to CO2 in the first hour. However, the conversion rate can easily be enhanced by finely tuning the reaction parameters, such as hydrogen peroxide concentration, temperature, and UV light intensity. The results of a run that was performed at 45⬚C and a relatively low initial phenol concentration show

Fig. 7. Total organic carbon concentration profile for several runs at different phenol initial concentrations and temperatures.

almost 20 wt.% organic carbon conversion to CO2 . It took approximately 4 h of continuous irradiation in the presence of H2 O2 to achieve 70 wt.% conversion of the organic carbons to inorganic carbons ŽCO2 .. The rest of the organic carbons were mainly found in the form of formic and oxalic acids.

5. Kinetic model For our kinetic study, a simplified reaction pathway for phenol oxidation by UV radiation combined with hydrogen peroxide is depicted in Fig. 8. In this model, all carboxylic acids are lumped together as ‘acids’. Furthermore, because the predominant acid was oxalic acid, we used its molecular weight as an average to estimate the molar concentration of the acids. In addition, we neglected Reaction 3 in the hydrogen peroxide photolysis mechanism ŽTable 2., because the reported rate constant is very low. We assumed that hydrogen peroxide concentration remains constant since it was very much in excess. H2 O2 decomposes by reaction 1 but forms by reaction 4 ŽFig. 8.. This assumption was verified by the negligible change in the hydrogen peroxide peak area Žconcentration . as well. In the kinetics, we assumed second order rates involving the organic species concentration and hydrogen peroxide concentration ŽFig. 5.. Experimental results showed that phenol oxidation rates depended on hydrogen peroxide concentration and the initial phenol concentration. There was an optimum hydrogen peroxide concentration where the rate was maximum. De Laat et al. Ž1994. in their UVrH 2 O 2 oxidation study of chloroethanes in dilute-aqueous solution reported a similar observation. Nicole et al. Ž1990. calculated the reactor wall reflection factor, which varied from unity for non-reflecting surfaces to 4.8 for aluminum-foiled quarts reactor walls. In our study, the reactor had

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Fig. 8. A simplified reaction pathway for phenol oxidation by UVrH2 O2 .

non-reflecting walls ŽPyrex-glass reactor. that had a reflection factor of unity. The estimated total absorbance of the solution was very high Ž At s 15.7., hence in Eq. Ž3., the term Ž1 y eyAt . and the equation was reduced to w OH . x ss s

2 ⌽ H 2 O 2 I0

Ž6.

n

k 2 C H 2 O 2 q k 5 C . OOH q

Ý

k i Ci

is1

Initial concetrations of hydrogen peroxide and phenol were known, and values of Io and l were measured and calculated ŽTable 1.. Based on Baxendale and Wilson Ž1956., the primary quantum yield of hydrogen peroxide Ž ⌽H 2 O 2 . is 0.49. In the overall process ŽReactions 1᎐4., two hydroxyl radicals are formed per photon absorbed. Hence, the overall value of ⌽H 2 O 2 was taken as unity. We used the literature values for the reaction constants k2 , k4 , and k5 ŽTable 2.. Initial concentrations of the hydroxyl radical were assumed to be zero. The term Ž k 5 C . OOH . in the denominator of Eq. Ž6. is negligible compared to the other two terms. Parameter estimation of the variables Ž ki . is optimized, based on how well the model fits a set of experimental observations. SimuSolv TM uses a systematic and quantitative method to adjust and optimize ki values. It uses the log likelihood functions as the criterion and either the generalized reduced gradient method or the Nelder᎐Mead search method to adjust the ki . Values of the hydroxyl radical concentration were substituted into the above equations and concentrations of phenol, catechol, hydroquinone, resorcinol, benzoquinone, acid, and CO2 were calculated at time t.

241

tion, the reaction rate constant of phenol disappearance Ž k6 q k7 q k8 . f 1.41= 1010 My1 sy1 agrees well with Apak and Hugul Ž1996. who reported rate constants of phenol photolysis as 1.4= 1010 My1 sy1. However, as the phenol concentration increases, the hydroxyl radical concentration decreases because phenol molecules compete more efficiently with hydrogen peroxide molecules for the radicals. Doubling the initial phenol concentration from 5.17= 10y3 to 1.05= 10y2 M decreased the hydroxyl radical concentration from 1.6= 10y11 to 5.5= 10y12 M. Hence, the photoassisted oxidation rate of phenol decreases. In addition, this suggested that the generated intermediates and the acids become increasingly important scavengers of hydroxyl radicals. As mentioned before, increasing hydrogen peroxide concentration enhances the oxidation rate up to an optimum H2 O2rphenol molar ratio where hydrogen peroxide starts to inhibit the degradation process. Therefore, strictly speaking, the hydrogen peroxide dosage does not significantly affect the reaction rate as long as the reaction is performed within the optimized molar ratio. Correspondingly, it is conceivable that the kinetics of photoxidation of organics depend mainly on the initial concentration of the parent organics. A plot of lnŽ k . vs. lnŽ CPh . demonstrates the inhibitory effect of the initial phenol concentration on the oxidation rates ŽFig. 9.. This negative dependence of the reaction rate on initial phenol concentration is correlated in the form of Eq. Ž7.. n . ln Ž ki . s ln Ž ai C Ph 0

Ž7.

The above discussion is in agreement with Augugliaro et al. Ž1988. findings that the pseudo-first-order reaction rate constant is inversely proportional to the initial phenol concentration. In Eq. Ž7., the slope n f y1 ŽFig. 9.. The linear relationship was incorporated with Eq. Ž6. to give the following ordinary differential

5.1. Initial phenol and hydrogen peroxide concentration effects For a given H2 O2 concentration range Ž; 50᎐500 mM., the degradation rate of phenol depends on its initial concentration. At low initial phenol concentra-

Fig. 9. Effect of initial phenol concentration on reaction rate constants of phenol oxidation.

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

242

equation network for the kinetic scheme shown in Fig. 8:

Table 3 Reaction rate constants

COH ⭈ dCPh Ž k6 q k7 q k8 . CPh s dt CPho

k i Žsy1 .

Temperature Ž o C.

k6 k7 k8 k9 k 10 k 11 k 12 k 13

27 6.51= 10 6 8.33= 10 3 1.27= 10 6 3.31= 10 6 3.83= 10 4 2.08= 10 6 1.00= 10 5 2.23= 10 4

Ž8.

COH ⭈ dCCatechol Ž k6 CPh y k9 Ccatechol . s dt CPho

Ž9.

COH ⭈ dCResorcinol Ž k7 CPh y k10 CResorcinol . s dt CPho

Ž 10 .

COH ⭈ dCHydroqunione Ž k8 CPh y k11CHydroquinone . s dt CPho Ž 11. COH ⭈ dCBenzoquinone Ž k11CHydroquinone s dt CPho yk12 CBenzoquinone .

Ž 12 .

COH ⭈ dCAcid Ž k9 CCatechol qk10 CResorcinol s dt CPho qk12 CBenzoquinone y k13 CAcid .

Ž 13.

COH ⭈ dCCO2 s k C dt CPho 13 Acid

Ž 14.

The above equations were coupled with the equation for hydroxyl radical concentration, Eq. Ž15.. COH .s s s

¡

~

2 ⌽ H 2 O 2 I0

k C ¢qk C 2

¥

§

w C Ž k q k7 q k8 . CPho Ph 6 Catechol q k10 CResorcinol k11CHydroquinone qk12 CBenzoquinone q k13 CAcidx

H 2 O2 q

9

1

¦

Ž 15.

5.2. Kinetic simulation We solved the above differential equations Eqs. Ž8. ᎐ Ž15. by numerical simulations using the SimuSolv TM software package. The simulation started by substituting the hydroxyl radicals steady-state concentration Ž COH . . in Eqs. Ž8. ᎐ Ž14.. The initial guesses for ki values were obtained from the intercept of the plot of

45 1.05= 10 7 2.24= 10 5 2.76= 10 6 9.53= 10 6 1.14= 10 7 5.44= 10 6 4.07= 10 7 1.00= 10 6

rate constants vs. initial phenol concentration, Eq. Ž7., and Fig. 9. A suitable solution for reaction rate constants Ži.e. k6 through k13 . was found based on the goodness of fit and the fitted rate constants are presented in Table 3. The predictions of phenol, hydroquinones, benzoquinones, acids, and CO2 as a function of time are shown in Figs. 10 and 11 by solid curves. Phenol degradation and hydroquinones production Že.g. catechol, resorcinol, and hydroquinone. were measured under many different conditions Že.g. temperature, H2 O2rphenol molar ratio, initial phenol concentration, etc... The predictions are excellent at most operating conditions and detailed results and analysis can be found in Alnaizy Ž1999.. The exception was the hightemperature, low initial phenol concentration run. The model over-predicted the reaction rate. To interpret this, it is important to recall that the reaction at low initial phenol concentration proceeds very fast. The phenol concentrations predicted by the model were lower than the observed data. At this relatively low phenol concentration, the estimated steady-state hydroxyl radical concentration is relatively high Ž; 10y11 M., but may not actually be as high as predicted by the model. In other words, the model over-predicted hydroxyl radical concentration, which in turn resulted in a faster oxidation rate. Predictions of the acids concentration were reasonably good for most reaction conditions; any deviation from the data was due to lumping all acids together since we could not identify all the acid products and hence could not calibrate the HPLC analysis. The predicted concentration of CO2 was significantly lower than the observed data. The observed values were at least twice those calculated. This discrepancy arises because CO2 concentrations were not determined very accurately. They were implicitly calculated as the difference between the total organic carbon reading from the TOC analyzer and the theoretical amount of organic carbon that was present in the phenol initially. Some of the error is also due to residual carbon dioxide

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Fig. 10. Observed and predicted concentrations of aromatic intermediates for phenol UVrH2 O2 oxidation, experimental conditions: T s 27⬚C; C0Ph s 1.07= 10y3 M; CH 2 O2 s 0.049 M; H2 O2rphenol molar ratio s 45.

in the solutions analyzed, giving artificially high experimental values.

6. Conclusions Results indicate that the oxidation rate was exclusively due to hydroxyl radical attack when hydrogen peroxide was in excess. It was shown that in the presence of hydrogen peroxide, the photodegradation process significantly increased with respect to UV radiation alone. At initial phenol concentrations of approximately 200 ppm, phenol completely disappeared in less than 1 h at 27⬚C by the UVrH2 O2 process, however, only approximately 20% of phenol was mineralized to CO2 and water and the rest was converted to the reaction intermediates. There was an optimum oxidantto-organics ratio for the photoxidation of organic pollutants. The optimal values and corresponding molar ratios of hydrogen peroxide to contaminant greatly affected the oxidation rates. In addition, results show

Fig. 11. Measured and predicted concentrations of acids and 0 CO2 . Experimental conditions: T s 27⬚C; CPh s 1.74= 10y3 M; CH2 O2 s 0.147 M; H2 O2 rphenol ratio s 84.

243

that as the initial contaminant concentration increased, the efficiency of the UVrH2 O2 AOP decreased. A free-radical mechanism involving hydroxyl radical reactions with phenol was developed. The first reaction step is the hydroxylation of the aromatic ring to yield hydroquinones, which further oxidize by hydrogen abstraction to yield benzoquinones. Then, the benzoquinone ring is cleaved to form muconic acids, which decompose by . OH to form maleic, fumaric, and oxalic acids. All intermediates formed initially are finally oxidized to mainly oxalic and formic acids which are finally destroyed after prolonged oxidation time by conversion to water and CO2 . Nonetheless, the reaction time is rather long, more than 2 h, at high initial phenol concentration Ž) 150 ppm.. A kinetic model for the advanced oxidation process using UV radiation combined with hydrogen peroxide was developed. Unlike most other kinetic models of UVrH2 O2 AOP, the model does not assume pseudo-first-order kinetics. The model provides a comprehensive understanding of the negative impact of initial phenol concentration on the photoxidation rate. The apparent rate constants were on the order of 107 sy1 at a temperature of 45⬚C and 106 sy1 at 27⬚C. References Apak, R., Hugul, M., 1996. Photoxidation of some mono-, di-, and trichlorophenols in aqueous solution by hydrogen peroxiderUV combinations. J. Chem. Tech. Biotechnol. 67, 221᎐226. Alnaizy R., 1999. Mechanism and Kinetic Study for the Hydroxyl Free Radical Mediated Photooxidation of OrganicContaminated Aqueous Solutions. PhD Dissertation. Texas A& M University. Augugliaro, V., Palmisano, L., Sclafani, A., Minero, C., Pelizzetti, E., 1988. Photocatalytic degradation of phenol in aqueous titanium dioxide dispersions. Toxicol. Environ. Chem. 16, 89᎐90. Baxendale, J.H., Wilson, J.A., 1956. The photolysis of hydrogen peroxide at high light intensities. J. Trans. Faraday Soc. 53, 344᎐356. Beltran, F.J., Ovejero, G., Rivas, J., 1996. Oxidation of polynuclear aromatic hydrocarbon in water. 3. UV radiation combined with hydrogen peroxide. Ind. Eng. Chem. Res. 35, 883᎐889. Buxton, G.V., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical review of data constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solutions. J. Phys. Chem. Ref. Data 17, 513᎐586. Christensen, H., Sehested, K., Corfitzen, H., 1982. Reactions of hydroxyl radicals with hydrogen peroxide. J. Phys. Chem. 86, 1588᎐1590. De Laat, J., Tace, E., Dore, M., 1994. Degradation of chloroethanes in dilute aqueous solution by HO2rUV. Water Res. 28, 2507᎐22519. Devlin, H.R., Harris, I.J., 1984. Mechanism of the oxidation of aqueous phenol with dissolved oxygen. Ind. Eng. Chem. Fundam. 23, 387᎐392.

244

R. Alnaizy, A. Akgerman r Ad¨ ances in En¨ ironmental Research 4 (2000) 233᎐244

Dulin, D., Drossman, H., Mill, T., 1986. The products and quantum yields for photolysis of chloroaromatics in water. Environ. Sci. Technol. 20, 72᎐77. Elliot, A.J., Buxton, G.V., 1992. Temperature dependence of the reaction OH q O2 and OH q HO2 in water up to 200⬚C. J. Chem. Soc. Faraday Trans. 88, 2465᎐2470. Glaze, W.H., Kang, J., Lay, Y., 1995. Advanced oxidation processes. A kinetic model for the oxidation of 1,2-dibromo-3-chloropropane in water by the combination of hydrogen peroxide and UV radiation. Ind. Eng. Chem. Res. 34, 2314᎐2323. Hong, A., Zappi, M.E., Kuo, C.H., Hill, D., 1966. Modeling kinetics of illuminated and dark advanced oxidation processes. J. Environ. Eng. 5, 58᎐62. La Grega, M.D., Buckingham, P.L., Evans, J.C., 1994. Hazardous Waste Management. McGraw-Hill, New York, NY. Lipczynska-Kochany, E., 1993. Hydrogen peroxide mediated photodegradation of phenol as studied by a flash photolysisrHPLC technique. Environ. Pollut. 61, 147᎐152. Masten, S.J., Davies, S.H.R., 1994. The use of ozonation to degrade organic contaminants in wastewater. Environ. Sci. Technol. 28, 180A᎐185A.

Nicole, I., De Laat, J., Dore, M., Duguet, J.P., Bonnel, C., 1990. Use of UV radiation in water treatment: measurement of photonic flux by hydrogen peroxide actinometry. Water Res. 24, 157᎐168. Ogata, Y., Tomizawa, K., Furuta, K., 1983. The Chemistry of Peroxides. John Wiley, Chichester, UK. Scheck, C.K., Frimmel, F.H., 1995. Degradation of phenol and salicylic acid by ultraviolet radiationrhydrogen peroxideroxygen. Water Res. 29, 2346᎐2352. Stefan, M.I., Hoy, A.R., Bolton, J.R., 1996. Kinetics and mechanism of the degradation and mineralization of acetone in dilute aqueous solution sensitized by the UV photolysis of hydrogen peroxide. Environ. Sci. Technol. 30, 2382᎐2390. Wei, T.Y., Wan, C.C., 1991. Heterogeneous photocatalytic oxidation of phenol with titanium dioxide powders. Ind. Eng. Chem. Res. 30, 1293᎐1300. Weinstein, J., Bielski, B.H., 1979. Kinetics of the interaction of HO2 and Oy radicals with hydrogen peroxide. The 2 Haber᎐Weiss reaction. J. Am. Chem. Soc. 110, 58᎐62.