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Environ. Scl. Technol. 1982, 16 676-681Printing Office No. 040-000-00337-1.U.S. nvironmental P rotection Agency-Region 11, DraftEnvironmental Impact S tatement for the Hudson RiverPCB Dem onstration Reclamation Projec t, May 1981.Fed. Regis t . 1976, 41, 7552.Choi, P. S.K.; Nack, H.; Flinn, J. E. Bull. Environ. Contam.Toxicol. 1974, 1 1 , 12-7.Law, L. M.; Goerlitz, D. F. Pestic. Monitor. J 1974,8,33-6.Carey,k .; Gower,J.A.; Tai,H.; Mitchell, W. G.; Wiersma,G. B. Pestic. Monitor. J . 1979, 12, 209-29.Carey, A. E.; Douglas,P.; Tai,H.; Mitchell, W. G.; Wiersma,

G . B. Pestic. Monitor. J 1979, 13, 17-22.Received for review November 11,1981. Accepted M ay 19,1 982.LDGO C ontribution No . 3356. Suppo rt for this research wassupplied by National Oceanic and Atmospheric Adminis tratio n(Contract NA79 -RAC OOl26 ) and the National Science Foun-dation (Con tract OCE-7909249). Add itional support was pro-vided by the New York State Department of EnvironmentalConservation (Contract NYS-C -12563 8) to implement Section3 of an agreement between the DEC and the General ElectricCo. related to the discharge of PCBs into the Hudson River.

Decomposition of Ozone in Water: Rate of Init iation by Hydroxide Io ns andHydrogen PeroxideJohannes Staehelln and Jurg HolgnQFeder al Inst i tu te for Water Reso urc es and Water Pol lution Control (EAWA G), Swiss Federal Inst i tu te of Technolog y,CH-8600 Dubendorf , Switzer land

W Th e initiation of ozone decomposition in pure wateris first order in O3 and OH- concentration and k03,0H-becomes 70 7 M-l s-l when a sequence of reactions isassumed by w hich three molecules of O3are eliminated perprimary event. Th e difference between this value andhigher values reported in the earlier literature may beexplained by interfe rences by radica l chain reactions whichhave not been totally inhibited in those studies. H 20 2 lsoreacts with O3when present as an anion, H02-. k03,HOz-is 2.8 0.5) X l o6 M-l s-l when it is assumed th at twomolecules of O3 are eliminated per primary event.Therefore, whenever [H,02] > lo- M (pH

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PH

02 40 -a * 100

Flgure 1. Scheme of reactions of aqueous ozon e. M solutes thatconsume ozone and becom e oxidized to MoxUand H,O,); I solutesthat initiate the decomposition of ozone to primary radicals; S, freeradical scavenger; M', solutes that reac t with OH. and form secondaryradicals R e and become finally oxidized to Mroxu.previously (3). Phos phate buffer solutions were preparedfrom phosphoric acid and sodium hydroxide (both su-prapu re grade). Th e solutions were preozonized with 3mg/L 03. odium bicarbonate was analytical grade. Itssolutions were preozonized. Nitrou s acid and sodiumhydroxide stan dard solutions were Merck Titrisol typereagents (0.1 N). HzOZ solutions were prepared fromNazOz (16)or from commercial perhydrol solutions thatwere free from stabilizers. Both reag ents gave th e sameresults. Th e H2 02 oncentrations were determined by themethod of Eisenberg 17).For methylmercury preparationsee ref 18.Instrumentation, pH measurements were performedwith a glass electrode. Th e calibration was based on MerckBuffer Titrisols for pH 7 and 10. Th e instrumental pHscale was calibrated for the determination of OH- con-centrations by solutions of defined OH- concentration(standa rd sodium hydroxides an d nitrous acid solutions)a t p H 10-11 (Gran titration method) 19). The ionicstrength for these calibrations was adjusted by N aN 03 tocorrespond to that used for the experiments on the kineticsof the OH- reactions I = 0.15 M).UV absorption measurements were performed on aBeckman DK-2A or Uvikon 810 instrum ent in thermo-stated and closed cells.Methods. Procedure for Kinetic Measurements.All experiments were performed a t 20f 1 C.Thermo-stating was required for measurem ents lasting more tha na short time. Samples from the concentrated aqueousstock solutio n of O3were pipetted to the prepared m ixturesof reactan ts. T he flasks were closed an d imm ediatelystirred or shaken for 5 s. For kinetic measurements per-formed to test th e effect of carbonate at pH 9 and for allseries with H 20 2, he depletion of O3vs. time was directlyfollowed in a 5-cm UV cell a t 258 nm. In a ll other casesth e ozonation was performed in a 1-L round bo ttom flaskfrom which series of 10-mL aliquots were withdrawn a tspecified tim e intervals. T h e residual concentration of O3in these samples was determined immediately by directUV measurem ents in 1-cm cells (for series at pH 8) or bythe indigo method (for all other series) 2~421).Procedure for Measurements of O H - RadicalYields. Th e yields of OH. radicals were determined fromthe stoichiometric factor by which methylmercury wasmineralized 18). This m ineralization is not disturbed byother ozonation reactions. For experimental details seeref 22.

p r i m a r yr a d i c a l s* 0 2 ,OH *

t ~ l e c o n d a r y t o2r e a c t i v e l o x idr a d i c a l s R

Flgure 2. Successive half-lives of ozon e vs. conc entration of carbo -nate for different pH values. [O3 I0= 50 p M : 1st fl,, A; 2nd t , , ,o. o3lO 3 p~ X; [ o 3 I 0= 0.3 PM +; [Polltot = 50 mM.

0.2 i \2A -0 20 40 60 80t i m e s

Flgure 3. Relative concentration of ozone vs. time for different [O&at high scavenger concentration: [HC03-] + [CO,*-] = 10 mM;[PO,], = 50 mM; pH 10.0.Results and Discussion

Decomposition Initiated by OH- Ions. Dependenceon pH and Scavenger and Ozone Concentration. Theconcentration of the O3 added could easily be followed asa function of time in the batch-type reactors when the pHwas in th e range 8-10. At pH values below 8, measure-men ts were still possible, but th e rates of elimination ofO3 became very small and could be disturbed by otherreactions suc h aswall effects. A t pH values above 10 thedecomposition of O3 became so fast th at i t could not befollowed accurately by th e ex perimen tal techniques appliedhere.From th e plots of O3concentration vs. time, half-livesfor the concentration of O 3 have been determined forsucceeding time intervals. As shown in Figure 2, thesehalf-lives depended m ainly on pH. At constant p H theyincreased when carbonate was added to scavenge OH.radicals (5). At low carbonate concentrations they alsovaried with initial O3concentration and time. However,when sufficient carbonate was added, the half-lives ap-proached a plateau value for each pH , and the y becameindependent of the O3 concentration. Similar results wereobtained when methy lmercury was used as a scavenger forOH. radicals (18).

Semi-log plots of O3 concentrations vs. time yieldstraight lines for those experiments performed a t highEnviron. Scl. Technol., Vol. 16, No. 10, 1982 677

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10-2

10-4J16-6 16-5 16 4[OH-] , M

a 9 10 -PHFigure 4. Measured pseudo-first-order rate constant for the decom-position of ozone vs. hydroxide lon concentration. A pH scale is givenfor comparison. [HCO,-1 + [C0,2-] = 10 mM; [PO,], = 50 mM.scavenger concentration as demo nstrated in Figu re 3 fora series of measu reme nts at pH 10. The slopes of theselines do not vary with O3concentration in the range 0.3-50KM. Therefore, the results presented in Figures 2 and 3for high concentrations of radical scavengers can be de-scribed by a pseudo-first-order ra te law:

- (d [O , ] /d t )p~ = kbH-[O$llO 1)where kbH-= In 2 / t l l z . As shown in Figure 4 the pseu-do-first-order rate constants t be increase by a facto r of10 per increased pH unit. For later convenience the pHscale in Figure 4 has also been conve rted to a scale ofincreasing OH- concen tration by the calibration proceduredescribed in the E xperim entalPart.T he slop e of log kbH-w. log [OH-] becomes 0.99 0.05 (error range for p = 0.95)when th e regression is based on seven series of m easure-ments performed at five different pH values. Thus, themea sured ra te of d ecomp osition of O3can be d escribed by

2)3)

d[O3]/ d t = ko,,o~-~~~[031~~[OH-]~~for which th e se cond-o rder rate-c onsta nt becomeskoa,OH-measd koH-/[OH-]l.O = (210 20) M-l s-lwhere [OH-] is the analytical OH- concentration. Whenbased on the experimental pH scale, the values becomesome what different because [OH-] is not exactly X10PHM 19).Transformed to the pH scale, the numericalvalues becomekog,pHmeasd= k b H - 1 0 1 ~ ~ (175 10)x 1014-PHM- s-1

3)Ranges of errors are base d on th e precision of th e mea-surements and on p = 0.95; th e accuracy of the dat a is,however, limited by th e accuracy of pH or [OH-] mea-sure me nts, for which we assu me a range of error of 0.05units or 10%.Discussion. On the basis of these results we assumetha t th e pH -dependent decomposition of O s is due to areaction tha t is kinetically controlled by the OH- concen-tration. The phenomeno n cannot be due to a pH-de-pendent change of O3 molecules (such as formation ofH 0 3 +or O,OH-) because all other known O3 eactions arestrictly pH indepen dent except in those cases where thedegree of dissociation of the other reactant changes withp H 21).Also the fact th at OH--initiated decompositionof O3 s of first order, both in Os and in OH- concentration,is consistent with our experience with reactions of O3withother solutes 21).

Th e kinetic rate laws now found for aqueo us solutionsthat contain OH. radical scavengers are different from678 Environ. Scl. Techno l., Vol. 16, No. 10, 1982

those reported in t he literature for pure water. Even theorders of reaction are different. This result is not sur-prising since in most earlier studies the rate of th e de-composition in pur e water was accelerated by undefinedchain reactions subsequent to OH- attack.The rate cons tant kOs,-,H-measd is small when compare dwith the rate c onsta nts quoted for other reactions of 03.This measured value is still different from the rate con-stan t of the prim ary OH- attack on O3 molecules becauseit must be assumed tha t more than one O3 molecule isdestroyed per primary initiation step. To get an estimationof the overall stoichiometry of the O3 + OH- reaction(when all OH. radicals are scavenged before reacting w ith0, he following reaction sequences can be considered:Hypothesis I ((2)exte nde d for dissociation of H Oz-andelectron-transfer mechanism)(4)

where(5)(6)(7)

HOz. H+ + - 0 2 - pK, = 4.8O3 e o z - s so3-+ O2

.03- H 2 0 H . + OH- + 0 2kOgOKprimO3+ OH- ___ HO2- + 02

H202 HOz- + H+ pK, = 11.6

Hypothesis I1 (this study an d ref 11)(4)

8)where

kOaHOZ-@O3+ HOz- H. + -0, O2followed by reactions 6 and 7.Hypothesis I11 (23)

(4 )followed by reactio n 7.Th e occu rrence of s o 2 - as an intermediate (hypothesesI and 11) is supported by q ualitative chemical detectionsbased on the reduction of tetranitrom ethane 22). T hetran sie nt ex istence of 03- as now been positively provedby kinetic spectroscopy (1 5 )as w ell asby pulse radiolysisstudie s on ozo nated water (kinetic spectroscopy) 24) .Hypothesis I11 was proposed for observations mad e a textreme conditions of very high pH values (pH -14). Forour system th e stoichiometric yield of OH. radical for-mation was found to be 0.55 f 0.08 (3). This value issignificantly less than 2.0 predicted by hypothesis 111, butit is close to 0.67 as predicted by hypotheses I and 11. Wetherefore can reject hypothesis I11 for our case.Because the observed overall reaction has bee n foundto be first order in O3 concentration, th e sequen ces ofreactions form ulated by hypotheses I and I1 must both beassumed to be rate limited only by the initiation step(reaction 4 or 4). This is expected the subsequent re-acti on of -0, ith O3 s fast and dom inant whenever theconcentration of O3 is dominant compared with that ofother solutes reactive toward so2-. t yields 03- 1624)(the rate con stant for the reaction of -0, ith 03, 6 sabout lo8M-I s-l ( 1 5 , 2 4 , 2 5 ) ) .Th is is relatively high whencomp ared with other r ate co nstants of -0, 26)).Likewise,SO,-nd H 20 2 an only exist as a short-lived intermediatesin the pH ranges here considered (see ref 24 and laterparagraphs).

All sequence s of reactio ns proposed by hypoth eses 1-111can be followed by OH. radicals interacting with f urthe r

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O3as form ulated by reaction 10. However, scavengers suchas carbonate, phosphate, or other solutes compete withreaction 10.

+HCO;OH.- H- + HC03. 12)+BP0,2- (13)H-- H- + .HP04- \ --,kI3 = 5 X lo6 M-l s-l 29)

In contrast to OH., H02. -02-)s not scavenged andleads to furth er decomposition of O3 (see eq 6).On th e basis of previous work 5 ) we assume that theproducts formed from bicarbonate and carbonate ions donot interact significantly with further 03, nd we assumethat reaction 6 is not inhibited by carbonate (reactionsbetween carbonate ions and H 02 - -0L) ave also not beenobserved in pulse radiolysis studies). Therefor e, for eachpH and each O3concentration an approp riate carbonateconcentration could be achieved that was sufficient toinhibit the subsequent chain reactions effectively.In both hypothesis I and I1 i t is assumed that threemolec ules of O3are decomposed per initiating ste p pro-vided that eo2-, 02., so3-,nd H202 re no t scavenged byother reactions. Herewith the rate constant of the primaryreaction of OH- with O3 becomeskOa,OHJrim = 1/33k0a,0H-measd = 70 f 7 M-l s- (14)

Th is value is in good agreement with th e value recentlydeduced by Forni et al. 15) for reactions a t pH >lo.Other Possible Primary Reactions in Pure Water.(i). A priori we cann ot exclude tha t th e decomposition ofaqueous O3could alsobe initiated by other species presentin pure water such a s by H30 +. H owever, the decompo-sition of O3 s not accelerated even a t very low pH values.At pH 2 tlI2s larger than 200000 s. Therefore k ~ ~ , H , o + ~ ~mu st be smaller than 4 X(ii). Aqueous O3could also decompose in th e aqueousmatrix by a spo ntaneous reaction tha t is not catalyzed byions. Th e rate law of this reaction could be formulatedby

-d [Od /d t = k [ o d (15)On th e basis the stability of O3a t low pH values, we canestimate that the first-order reaction rate constant for sucha reaction, tz,, is smaller tha n 5 X lo4 s-l. Thu s, initiationof O3 decomposition reactions that are not controlled byOH-, such as proposed by Alder and Hill (30 ) , annot beof importance.Decomposition Initiated by H 2 0 2 H 0 2 - Ions).Measurements on th e ra te of the decomposition of O3 nthe presence of radical scavengers were performed on aseries of samples th at contained different concentrationsof H 2 0 2 . For these measurements methylmercury hy-droxide was used as a free radical scavenger instead ofcarbonate 18) (carbonate would be inefficient in th e lowpH region applied, where it is protonated to form H2C0,).H202 was added in an excess when comp ared with t ha t ofth e initial concentration of 03.All plots of the log of th econcentration of residual O3vs. time were linear. Th us ,pseudo-first-order rate constants for th e elimination of O3

M -l s-l.

Flgure 5. Measured pseudo-flrst-order ra te co nstants of decompositionof ozo ne vs. hydrogen peroxlde concentration for different pH values:[methylmercury] = 1.5 mM; [PO4ltOt 50 m M.

0 2[ H O I ,lo- M

Figure 6. Measured pseud o-flrst-order ate constant of decompositionof ozone vs. hydrogen peroxlde ion conc entration: [methylmercury]= 1.5 mM; [P04],ot= 50 m M.could again be deduced from t he slope of these lines. A tpH 2, H 2 0 2 eacts only very slowly with O3 4) . However,at p H values above 5, a strong acceleration of t he decom -posi tion of O3 by H20 2was observed. The effect increasedlinearly with th e concentration of H 20 2 see Figure 5).Therefore the reaction rate can be formulated by

-(d[031/dt)p1-1 = ~ 0 ~ ~ 0 ~ ~ ~ ~ [ 0 3 1 ~16)This r ate increases by 1order of magnitude per pH unit.This observation suggests tha t H 20 2 lso reacts with O3only when present in its ionized form:

V e r y S l o w0 3 H20 2 - see r e f 403 + Hoe- d e c o m p o s i t i o n o f O 3

T he r ate of decom position of O3 s plotted in Figure 6 vs.th e con centration of H02-. Th e concentration of H 02 -wascalculated from the total concentration of H 2 0 2 nd thedegree of dissociation, a, alculated for pKH&, = 11.6 andthe p H applied, wherea = [HO,-I/([H20,1 + WO2-I) = [HO,-l/[H2021,, (17)

(18)or for pH

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Tab le I. Rat e Cons t an t s for t he H,O-0, a n d t h e A q u e ou s H ,O , -0 , Systems210 20 M-] s l

H,O k, = < 5 X s- kO,.H,O+= k o , , o ~ - ~ ~ ~ k'OH- = 1 7 5 x 1OPH-14 s 1a

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105 I \ I I\

IO6 6PHFigure 8. Half-ilves f,,*) of ozone in different waters: I a11 OH.radicals scavenged (this work); 11, distilled wat er, [PO,], = 50 mM;111, water of the lake of ZiSrich (30); V, groundwater, infiltrated fromRhlne (Basle, Waterwo rk Hard) (37).

Rea ctions of OH- nd H02- with O3 an initiate radicalchain reactions. Th e kinetic chain length of such reactionsdepends on the relative rate by which the ra dicals formedreact with O3 ompared with with other solutes presentin the solution. Therefore such chain reactions can bemore important in pure water than in real drinkingwaters or in some typ es of waste waters where bicarbonateand organic impurities may significantly scavenge OH.radicals. In Figure 8, th e half-lives of O3 measured indifferent types of natu ral waters are comp ared with thosemeasured in our model solutions. Th is comparison alsodemonstrates t ha t our conclusions are applicable to mea-suremen ts on real waters.Acknowledgments

We thank Werner Stumm for his continuous supportand interest in this study and Werner Haag, Ren6Schwarzenbach,and Herbert Hohl for v aluable discussionsand reviewing th e manuscript. Heinz Bader was of greathelp in providing many of th e experimental techniquesapplied for these studies.Literature Cited(1) Hoigng, J. In Handbook of Ozone Technology and itsApplication; Rice, R. G., Netzer, A., Eds.; Ann ArborScience: Ann Arbor, MI, 1982.(2)Weiss, J. Trans. Farad. SOC. 936, 31, 668-81.(3) Hoign6, J.; B ader, H. Water Res . 1976,10, 77-86.(4)Taube, H.; Bray, W. C. J. Am . Chem. SOC. 940,62,3357-73.

5) Hoigng, J.; Bader, H. Vom Wasser 1977, 8, 283-304.(6)Stumm , W. Helv. Chim. Acta 1954, 7, 773-8.(7) Peleg, M. Water Res . 1976, 10, 361-5.(8)Roth, J. S.; Sullivan, D. E.; DAmbrosi, M. OzonewsTechnical Pape r Section 1979, .(99 Gurol, M. D. Dissertation University of N orth Carolina,Chapel Hill, NC, 1980.(10)Staehelin,J.; Hoigng,J. 5th Ozon-Weltkongress,April 1981,Wasser, Berlin, Colloquium Verlag: West Berlin , 1981; p623-636.(11) oy, P.;Gilbert, E.; Eberle, S. H. Water Res . 1980, 4,1509-16.(12)Gilbert, E. Water Res . 1980, 14, 1637-43.(13)Kruithof, J. C. In Oxidationsverfahren in derTrinkwasseraufbereitung;Kiihn, W., Sontheimer, H., Eds.;DVGW Karlsruhe, 1978; 88-416.(14)Yamamoto, Y.; Niki, E.; Shiokawa, H.; Kamiya, Y. J. Org.Chem. 1979,44, 137-42.(15)Forni, L.;Bahnem ann, D.; Hart, E. J. Phys . Chem. 1982,

(16) Riesenfeld, E. H. Anorg. Chem. Praktikum , Rascher-Verlag,(17)Eisenberg, G. M. nd. Eng. Chem. 1943,15, 27-8.(18)Hoigng, J.; Bader, H. n Organometals and Organo-metalloids, Occurrence and Fate in the Environment;

Brinckmm, F.E., Bellama,J. M., Eds., American ChemicalSociety: Washington, D.C., 1978;ACS S ymp. Ser . 82,(19)Stumm, W.; Morgan, J. J. Aquatic Chemistry, 2nd ed.;(20)Bader, H.; Hoigng, J. Water Res . 1981, 15, 49-56.(21)Hoign6, J.; Bader, H. Water Res . , in press.(22) Staehelin, J. Thesis, ETH Ztirich, in preparation.(23)Gorbenko-Germanow;Koslova Dokl. Akad. Nauk SSS R(24)Btihler, R.; Staehelin, J.; Hoigng, J., unpublished data,ETH-Z and EAWAG, 1982.(25)Sehested, K.; Holeman, J.; Ha rt, E. J., to be submitted forpublication.(26)Farhataziz; Ross, A. B. Selected Specific Rates of Reactionsof Trans ients from Water in Aqueous Solution, National

Bureau of Standards, Washington D.C., NSRDS-NBS 59,1977.(27)Bahnemann, D.; H art, E. J. Phys. Chem. 1982,86,252-255.(28)Weeks,J. L.;Rabani, J. J . Phys . Chem. 1966, 70,2100-6.(29)Black, E. D.; Hayon, E. J. Phys. C hem. 1970,74,3199-203.(30)Alder, M. G.; Hill, G. R. J . Am. Chem. SOC. 950,72,1884-6.(31)Hoign6, J.; Bader, H. Ozone Sci. Eng. 1979, , 357-72.

86 55-259.Zurich, 15,Auflage 1940, 356-7.

292-313.Wiley: New York, 1981; p 226-9.

1973,210, 51-4.

Received for review December 23,1981.Accepted June 9,1982.

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