HAZARDOUS WASTE & HAZARDOUS MATERIALSVolume 4, Number 2, 1987Mary Ann Liebert, Inc., Publishers

Destruction of Benzene by UltravioletLight-Catalyzed Oxidation with

Hydrogen PeroxideB. A. WEIR, D. W. SUNDSTROM, and H. E. KLEI

Department ofChemical EngineeringThe University ofConnecticut

Storrs, CT 06268

ABSTRACT

The removal of low levels of hazardous organic compounds from water is oftendifficult and expensive by conventional treatment methods. A promising method inwhich pollutants are oxidized by ultraviolet (UV) light-catalyzed hydrogenperoxide has been investigated for an aromatic pollutant, benzene. The oxidationswere conducted in a quartz annular reactor equipped with a 254 nm germicidal UVlamp. Benzene was more rapidly destroyed by the combination of hydrogen peroxideand UV light than by UV light alone. The rate of oxidation could be increased byincreasing the concentration of hydrogen peroxide initially present in the reactoror by increasing the UV light intensity. Alkaline pH was detrimental to thereaction rate, probably because of the base-catalyzed decomposition of hydrogenperoxide. A rapid increase in the absorbance of treated benzene solutions at 254nm was observed, and attributed to the presence of strongly absorbing aromaticoxidation intermediates. Four intermediates were identified by GC/MS and HPLCanalyses. The intermediates could be removed by adding additional hydrogenperoxide and extending the treatment time.

INTRODUCTION AND BACKGROUND

The removal of low levels of hazardous organic compounds from water can be a

difficult and expensive task. Conventional treatment methods such as packed bedaeration and granular activated carbon adsorption can effectively remove some

compounds but not others. Aeration is only useful for highly volatile pollutantsand is not practical for small scale use such as by homeowners with polluted wells.Carbon adsorption has the disadvantage that the carbon must be replaced or

regenerated when its adsorptive capacity is reached.An alternative to aeration and adsorption is the oxidation of organic pollutants

in water by the combination of ultraviolet (UV) light and a chemical oxidant. UVlight accelerates the rate of removal of a pollutant by activating the oxidant andin some cases by rendering the pollutant more susceptible to oxidation.

A process in which UV light-catalyzed ozone acts as the oxidant has receivedconsiderable research attention. The process has been proven effective for theoxidation of numerous organic compounds in water. Among these are the pesticides

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pentachlorophenol, roalathion, Baygon, DDT, and Vapam (1); the explosive TNT (2);the halogenated aliphatics chloroform, bromodichloromethane, andtetrachloroethylene (3); and the aromatics nitrobenzene, benzoic acid,3-chlorobenzoic acid, and anisóle (4).

While ozone/UV treatment is undoubtedly effective on a wide range of compounds,it has a number of disadvantages. Ozone is an unstable gas and must be generatedon-site and used immediately. An ozone contacting device must be provided whichcan achieve adequate mass transfer of ozone into the liquid phase. Ozonegenerators and contactors are probably too expensive for small industrial or

domestic water treatment systems.An oxidant which may be as effective as ozone but is better suited for use in

small treatment systems is hydrogen peroxide. Since hydrogen peroxide is a

relatively stable liquid, it may be stored for later use. There is no problem ofmass transfer between phases so elaborate contacting devices are unnecessary.

The initial step in the oxidation of organic compounds by hydrogen peroxide andUV light is believed to be the absorption of a UV photon by a hydrogen peroxidemolecule, causing it to dissociate into two hydroxyl radicals. The hydroxylradical has an oxidation potential of 2. 87 volts relative to hydrogen, which makesit a stronger oxidizing agent than either ozone or hydrogen peroxide (5). Theradical attacks organic compounds by abstracting hydrogen atoms or by adding todouble bonds.

The hydrogen peroxide/UV process has been proven effective for the treatmentof TNT-containing wastewaters (6), the removal of total organic carbon from tapwater and boiler feedwater (7, 8), and the oxidation of acetate in shipboardwastewaters (9). A study of the common drinking water pollutant trichloroethylene(5) has shown that the hydrogen peroxide concentration has the greatest effect onthe rate of oxidation of this pollutant. A comparative study of the rates ofdestruction of several saturated and unsaturated halogenated aliphatic compoundshas shown that unsaturated compounds are in general degraded more rapidly (10).

The purpose of this research was to demonstrate the utility of hydrogenperoxide/UV treatment for the removal of benzene from water. Benzene is aconfirmed human carcinogen and is one of the Environmental Protection Agency'spriority pollutants. Process variables such as the initial hydrogen peroxideconcentration, the UV light intensity, the temperature, and the pH were studiedto assess their effects on the reaction rate of benzene. For some runs thedestruction of the aromatic ring was followed by HPLC analysis, GC/MS analysis,or monitoring the absorbance of the treated solution at 254 nm.

EXPERIMENTAL METHODS

Annular Reactor

The recirculating flow reactor system is shown in Figure 1. The annular reactorwas made entirely of quartz. It was 20 cm long and had inner and outer diametersof 2.5 cm and 5.4 cm respectively, yielding a total annular volume of 285 mL. Alow pressure ultraviolet germicidal lamp (American Ultraviolet Co., Model G10T51/2 L) of 1. 58 cm diameter was positioned inside the reactor. The lamp had an

output of approximately 5.3 watts at 254 nm, which varied somewhat withtemperature. For this reason a thermocouple was attached to the surface of thelamp and its temperature was kept constant by blowing a stream of air along itssurface.

The reactor was operated in a recirculating mode: solution was pumped throughthe reactor, into a 3000 mL reservoir, and back into the reactor. The reservoirwas jacketed for temperature control and contained ports for sampling, addingreagents, and a thermometer. The flow rate was maintained at 1.8 gpm to ensure

adequate mixing. Complete mixing and the small size of the reactor relative tothe reservoir allowed the reactor to be modeled as a batch reactor.

The reactor system was prepared for a run by filling it with 2800 mL of 0.05 Mphosphate buffer at the desired pH. A separate vessel was used to dissolve the

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benzene in 500 mL of buffer. For all runs the initial benzene concentration was

approximately 0. 2 mM. Upon dissolution the benzene solution was added to thereactor along with the proper amount of "^50% hydrogen peroxide (Fisher ScientificCo. , ACS Reagent grade). After allowing 15 minutes for mixing, the lamp was turnedon. Samples were taken at intervals and analyzed as soon as possible for benzeneand hydrogen peroxide. For some runs a tubing pump (Cole-Parmer Co.) was used topump a small amount of the solution continuously through a quartz flow cell of 1.0cm path length. The flow cell was positioned in a Perkin-Elmer Sigma 3spectrophotometer set at 254 nm. The spectrophotometer was connected to a

Perkin-Elmer strip chart recorder which provided a continuous trace of theabsorbance of the solution over time.

Analysis Methods

Samples were analyzed for benzene concentration by direct injection into a

Gow-Mac Model 69-750 gas Chromatograph equipped with a flame ionization detector.A 10 foot by 1/8 inch OD stainless steel column packed with 3% SP-1500 on 80/120mesh Carbopack B was used. Peak area on the chromatogram was converted to benzeneconcentration by reference to a calibration curve.

Samples were analyzed for hydrogen peroxide by a modification of the glucoseoxidase

-

peroxidase method for analyzing glucose (11). The peroxidase enzymeconverted hydrogen peroxide to water and at the same time oxidized a leuco dye,whose absorbance was a measure of the concentration of hydrogen peroxide. Detailsof the method are described elsewhere (12).

Samples were analyzed for reaction intermediates by high performance liquidchromatography (HPLC). Samples of 30 mL volume were acidified and extracted intoethyl acetate, which was then evaporated under a stream of nitrogen. The residuewas dissolved in 200 uL of methanol and chromatographed on a Radial-Pak C^g columnunder a linear solvent gradient from 0. 1% H3PO4 in water to 0. 1% H3PO4 inacetonitrile. A Waters modular HPLC system was used with UV absorbance detectionat 220 nm.

Samples were also subjected to GC/MS analysis in a search for reactionintermediates. The extraction procedure was the same as that used for the HPLCanalyses, except that a larger sample volume of 500 mL was extracted. TheHewlett-Packard GC/MS system consisted of a Model 5890A gas Chromatograph connectedto a Model 5970 mass selective detector. A 12. 5 m, 0.2 mm ID crosslinked dimethylsilicone capillary column (HP Part No. 19091-60312) was used with the followingtemperature program: hold at 50^C for 5 minutes, then heat to 250^C at S^C/min,then hold at 250°C for 20 minutes.

RESULTS AND DISCUSSION

Effect of Process Variables

Figure 2 illustrates the synergistic effect of hydrogen peroxide and UV lightfor the destruction of benzene. Treatment with hydrogen peroxide alone caused no

reduction in benzene concentration but UV light alone did have some effect,reducing the benzene concentration by about 50% after 90 minutes. Apparently theabsorption of 254 nm light by benzene can cause it to decompose. The combinationof hydrogen peroxide and UV light is much more effective than UV light alone; over

98% of the initial benzene was removed after 90 minutes of treatment. The benzeneis probably being attacked by both UV photons and by hydroxyl radicals generatedfrom the hydrogen peroxide.

The effect of the initial hydrogen peroxide to benzene molar ratio (R) on theoxidation rate of benzene is shown in Figure 3. As the ratio is increased, thereare more hydroxyl radicals available to attack the benzene ring and it is destroyedmore rapidly. The semi-log plots are linear for the first 10 to 20 minutes,suggesting that the rate law for benzene oxidation may be first order in benzene.

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Circulator^ ¿ Rotameter

Recorder

AirSupply

Rotameter

FIGURE 1. Schematic diagram of recirculating flow reactor.

ai. 0.02

60 80Minutes

FIGURE 2. Effect of H202 alone, UV alone, and H202 plus UV on

decomposition of benzene at 25 ^C , pH 6.8, maximumintensity. Initial benzene = 0. 2 mM. InitialH202/benzene = 6. 6 mols/mol.

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Etuer

0.5 r

0.2 r

2 0.05

0.02

0.0140 60 80

TIME, iVinutes100

FIGURE 3. Effect of initial molar ratio of H202 to benzene on

decomposition of benzene at 25 °C , pH 6.8, maximumintensity. Initial benzene = 0. 2 mM.

CT>c'coEor

co

o

20 40 60 80TIME, Minutes

FIGURE 4. Effect of initial molar ratio of H202 to benzene on

decomposition of H202 at 25 0c , pH 6.8, maximumintensity. Initial benzene = 0. 2 mM.

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The deviations from linearity at later times may result from the decrease ofhydrogen peroxide concentrations with time (Figure 4). Alternatively, thedeviations may be caused by a buildup of reaction intermediates which absorb UVlight and/or react with hydroxyl radicals.

Pseudo first order rate constants have been calculated from the initial linearportions of the curves and are shown in Table 1. Also presented are the ratiosof each benzene rate constant to the rate constant for the action of UV light alone.This ratio shows the degree to which the reaction rate is accelerated by theaddition of hydrogen peroxide. The value of the ratio for a given run is roughlyequivalent to the value of R for that run. A doubling of the R value wouldtherefore be expected to approximately double the first order rate constant.

TABLE 1

Pseudo First Order Rate Constants at Different Initial Molar Ratios (R) of H2O2to Benzene. 25^C, pH 6.8, Maximum Intensity. Initial Benzene Concentration = 0.2mM.

1 k for UV + H202Benzene k, min"1 —;—-——;-k for UV alone

0.0 0.0145 1.02.5 0.0518 3.63.2 0.0693 4.86.6 0.0996 6.910.6 0.1636 11.3

Experiments conducted at 15, 25, and 35^ C showed that the reaction rates ofboth benzene and hydrogen peroxide are insensitive to temperature, which is notunusual since photochemically initiated reactions often have low activationenergies (13).

The reaction rate of benzene was considerably slowed by an alkaline pH (Figure5), probably because the hydrogen peroxide decomposed so rapidly under the basicconditions (Figure 6). Hydrogen peroxide is known to undergo base-catalyzeddecomposition. One proposed mechanism for the decomposition involves theperhydroxyl (HO2 —) ion (14):

H202 + H02-

-1 H20 + 02 + OH-

This type of decomposition is undesirable because it consumes hydrogen peroxidewithout generating reactive free radical species.

The effect of varying the incident UV light intensity on the reaction rate ofbenzene is shown in Figure 7. As the intensity is reduced from its maximum value,the benzene is oxidized more slowly, since fewer UV photons are being supplied tothe reaction mixture per unit time. A similar trend was observed for the hydrogenperoxide reaction rate.

Reaction Intermediates

When benzene solutions were treated with hydrogen peroxide and UV light, therewas an initial rapid increase in the absorbances of the solutions at 254 nm (Figure8). At all three temperatures the absorbance reached a maximum after about 1 hourand was eventually brought to zero by continuing the run and adding hydrogen

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0.05

0.02

0.01

pH 10.5

40 60 80TIME , Minutes

FIGURE 5. Effect of pH on decomposition of benzene at 25 ^C ,

maximum intensity. Initial benzene =0.2 mM. InitialH202/benzene = 7 mols/mol.

E01a.

£ 0.05^

0.02

0.0120 40 60 80

TIME, Minutes

FIGURE 6. Effect of pH on decomposition of H202 at 25 °C ,

maximum intensity. Initial benzene = 0. 2 mM. InitialH202/benzene = 7 mols/mol.

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0.05 h

0.02

00140

TI ME60 80 100Minutes

FIGURE 7. Effect of relative UV light intensity on

decomposition of benzene at 25 °C , pH 6. 8. Initialbenzene = 0.2 mM. Initial H202/benzene = 7 mols/mol.

FIGURE 8. Change of absorbance at 254 nm with time for runs at15, 25, and 35 °C, pH 6.8, maximum intensity. Initialbenzene = 0.2 mM. Initial H202/benzene = 7 mols/mol.H202 added at half-hour intervals.

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peroxide at half-hour intervals. The rapid increase in absorbance is caused bythe formation of reaction intermediates which strongly absorb UV light at 254 nm.

The fact that the absorbance curves are similar suggests that the reaction ratesof the intermediates are insensitive to temperature.

The buildup of strongly absorbing reaction intermediates is at least partiallyresponsible for the observed slowdown in the reaction rates of benzene and hydrogenperoxide. Since the rates depend on the UV light intensity, the rates for bothcompounds decrease as intermediates consume more of the available UV light. Itis also probable that the intermediates compete with benzene for the availablehydroxyl radicals.

The subject of intermediates was further considered by conducting HPLC analysesof treated benzene samples, the results of which are shown in Figure 9. Phenol,catechol, hydroquinone, and resorcinol were tentatively identified by theirretention times. A number of other peaks were present which were not identified,but it was shown that these peaks could be eliminated by extending the treatmenttime.

GC/MS analyses of treated samples confirmed that phenol, catechol, andhydroquinone are reaction intermediates. No compounds other than these were

observed in the GC/MS scans. The smaller number of compounds observed by GC/MSas compared to HPLC may result from the difficulty of detecting highly polar,non-volatile intermediates, such as aliphatic acids, by the GC/MS procedure usedhere. Further work with GC/MS in which derivatization techniques are used isnecessary to determine whether acids are formed. Since the procedure should havebeen able to detect biphenyl or phenolic dimers, the absence of such compoundssuggests that dimerization is not a significant reaction at the low concentrationsof benzene used in this work. As in the case of the HPLC scans, phenol and thedihdroxybenzenes disappeared with longer reaction times.

Correlation of Results

An empirical rate expression has been used to correlate the benzene rate databecause rate laws derived from simplified reaction mechanisms were unable torepresent the data adequately. The expression used was a two term rate expressionsimilar to that proposed by Glaze et al. (3) for the ozone/UV process:

rB = kp CB + kB CB C§where rB = the reaction rate of benzene, umol/L-min

kp = the rate constant for benzene photolysiskB = the rate constant for benzene oxidation by both hydrogen

peroxide and UV lightCB = the benzene concentration, umol/LCjj = the hydrogen peroxide concentration, umol/L

and a, b, and c are reaction orders.

While this expression has no theoretical basis, a knowledge of concentrationdependence is useful for design purposes.

The first term in the expression represents the reaction rate due to UV lightalone, while the second term describes the increase in rate due to hydrogenperoxide. Multiple linear regression analysis to find values for the rateconstants and reaction orders yielded the following rate expression:

rB = 7.55 x 10"6 Cjf4 + 3. 04 x 10"4 C¿- ° C§- 82The curves predicted by this rate equation show fair agreement with the data

points (Figure 10). The rate expression could perhaps be improved by taking into

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aJkjutJH

3 12

I Hour

IWuJaJau-a^^-

FIGURE 9. Liquid Chromatograph traces for benzene samples atstart of run and after 1, 2, 3, and 4 hours oftreatment. Peak 1 = benzene. Peak 2 = phenol. Peak3 = catechol. Peak 4 = resorcinoi. Peak 5 =

hydroquinone.

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account the effects of the incident UV light intensity and the absorption of UVlight by reaction intermediates.

0.5

0.2 y

.2 0.05

0.02 h

0.020 40 60 80 100

TIME, Minutes

FIGURE 10. Comparison of experimental benzene concentrationswith values predicted by kinetic model. Parameteris initial molar ratio of H202 to benzene.

CONCLUSIONS

1. Hydrogen peroxide and UV light acted synergistically to oxidize low levelsof benzene in water.

2. The rate of benzene oxidation was increased by increasing either the hydrogenperoxide concentration or the UV light intensity.

3. Benzene was oxidized more slowly at alkaline pH, probably because the hydrogenperoxide was undergoing a base-catalyzed decomposition.

4. The effect of temperature on the reaction rate of benzene was minimal.5. The absorbance of treated solutions at 254 nm changed dramatically over time

due to the formation and destruction of stror.gly absorbing reactionintermediates.

6. HPLC and GC/MS analyses demonstrated that several reaction intermediates couldbe eliminated by adding hydrogen peroxide and extending the treatment time.

7. An empirical rate expression was developed to correlate the benzene rate data.

ACKNOWLEDGEMENTS

This research was supported by the U. S. Department of the Interior under grantnumbers 14-08-0001-G-832 and 14-08-0001-G-1064. We also thank Dr. Dennis Hill forhis many helpful suggestions on the analytical aspects of the research.

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REFERENCES

1. Mauk, CE. , H.W. Prengle, and J. E. Payne, Oxidation of Pesticides by Ozoneand Ultraviolet Light, Houston Research Inc., Houston, Texas, 1976.

2. Fochtman, E.G. and J. E. Huff, Proc. Second International Svmp. on OzoneTechnology. R. G. Rice, P. Pichet, and M. Vincent, editors, International OzoneInst. , 1976.

3. Peyton, G. R. , F. Y. Huang, J. L. Burleson, and W. H. Glaze, Environ. Sei.Technol.. 16. 448 (1982).

4. Leitis, E. , An Investigation into the Chemistry of the UV-Ozone PurificationProcess, Final report to NSF, Grant no. ENV 76-24652, February 8, 1980.

5. Nalette, T. A. , Ultraviolet Light Catalyzed Oxidation of Trichloroethylenewith Hydrogen Peroxide, M.S. thesis, University of Connecticut, 1982.

6. Andrews, CA., Photooxidative Treatment of TNT Contaminated Waste Waters,Report no. WQEC/C 80-137, Naval Weapons Support Center, 1980.

7. Malaiyandi, M. , M. Sadar, P. Lee, and R. O'Grady, Water Research. 14. 1131(1980).

8. Clarke, N. and G. Knowles, Effluent and Water Treatment Journal. 22. 335(1982).

9. Koubek, E. , Ind. Eng. Chem. Process Des. Develop. . 14. 348 (1975).

10. Sundstrom, D.W., H. E. Klei, T. A. Nalette, D. J. Reidy, and B. A. Weir, HazardousWaste and Hazardous Materials. 3. 101 (1986).

11. Boehringer Mannheim GmbH, Biochemica Catalogue. 1971 edition, Mannheim,Germany.

12. Weir, B. A. , The Oxidation of Benzene in Water by Means of Hydrogen Peroxideand Ultraviolet Light, M.S. thesis, University of Connecticut, 1985, p. 115.

13. Calvert, J. G. and J. N. Pitts, Jr. , Photochemistry. John Wiley and Sons, Inc. ,

New York, 1966, p. 646.

14. Schumb, W. C. , C.N. Satterfield, and R. L. Wentworth, Hydrogen Peroxide.Reinhold Publishing Corp., New York, 1955, p. 476.

Address reprint requests to:

Dr. Donald W. SundstromDepartment of Chemical Engineering

The University of Connecticut191 Auditorium Road, U-139

Storrs, CT 06268

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