255
J. Phys. Chem. 1982, 86, 255-259
HzO2,and CH,, the gas-phase rate is actually higher than the rate in water.'vz4 The exceptionally high rate of reaction 2 in water relative to the gas phase may indicate that a rapid addition reaction occurs in aqueous solutions. OH. + O3 HO,. (5) While we are unable to decide the course of this reaction, our results indicate that the radical formed in reaction 5 is unable to react further with the acetate ion to produce its radical ion with a rate constant high enough to affect our measurements. It is possible, however, the HO,. decomposes into H02. + O2 as observed in the gas phase. The dissociated form of H02., 02-.,is known to react with O3with a rate constant of about 1 X lo9 dm3mol-' s-' (ref
-
23) to form and subsequently an OH. radical. This sequence of reactions could explain the pronounced slow development of .CH2COO- after the initial rapid reaction (Figure 2b). 0 3 - 0
Acknowledgment. We thank Professor K.-D. Asmus for his encouragement in undertaking this project and for discussions of the results. We are also indebted to K. Sehested and J. Holcman of the Ris0 National Laboratory, Denmark, for valuable discussions on the reactions of ozone with the hydroxyl radical. The technical assistance of Frau E. Darnsti4dt is gratefully acknowledged. E. H. also thanks the Hahn-Meitner-Institut for use of their laboratory and the Alexander-von-Humboldt-Stiftungfor a Senior U.S. Scientist Award that made these studies possible.
Mechanism of the Hydroxide Ion Initiated Decomposition of Ozone in Aqueous Solution L. Fornl,+ D. Bahnemann, and Edwln J. Hart' Hehn-Meltner-Instltur fOf Kernforschung Berlln GmbH, Berelch Strahlenchemle, P 1000 West Berlln 39, Federal Republlc of Germany (Received: July 15, 1981; In Flnal Form: September 4, 1981)
Stopped-flow experiments are reported on the OH--catalyzed chain decomposition of ozone in the pH range 11-13. 03-. has been identified as a product by ita 430-nm absorption band. Acetate and carbonate ions inhibit this reaction. In the carbonate-inhibited reaction, the radical anion C03-. has been identified by ita 600-nm absorption band. In this case an apparent second-order rate constant of 115 f 40 dm3mol-' s-' has been obtained. Evidence is provided to support the reaction OH- + O3 HOf + O2as the primary step. On this basis, k(OH+ 0,)equals 48 f 12 dm3mol-' s-l. Some aspects of the overall reaction are discussed.
-
Introduction The chemistry, as well as the biochemistry, of ozone has been of major interest for many scientists over the last In water chemistry ozone is used as an environmentally satisfactory and potent oxidant for purification purpose^.^ Although highly reactive free radicals have been reported as intermediates in many ozone reactions, their detailed mode of action has not been revealed unambiguously." The hydroxide ion induced decay of ozone, in particular, has been the subject of a number of studies.7-ls Because of the rapidity and complexity of this reaction in alkaline solutions, an accurate determination of the ozone decomposition kinetics has been difficult to work out. The present investigation has been undertaken by using a time-resolved stopped-flow technique with a resolving time of about 1ms. Previous work has established the decomposition of ozone above pH 8 as a bimolecular process13J6with a rate given by -d[031/dt = ki([031[OH-l) (1) However, there is considerable uncertainty regarding the value of kl for the overall decay of ozone, typically quoted values being 700 (ref 15) and 1500 (ref 17) dm3 mol-l s-l. The decay of ozone in solution appears to be a chain process involving the hydroxyl radical,'-9 where the hyBrunel University, Department of Biochemistry, Uxbridge, Middlesex UB8 3PH, Great Britain. *Award of the Alexander von Humboldt-Stiftung, 1979-1980. Address correspondence to this author at the following address: 2115 Hart Road, Port Angeles, WA 98362. 0022-365418212086-0255$01.25/0
droxide ion is involved as an initiator. Thus, in reactions which involve ozone above a certain pH, the hydroxyl radical and other breakdown products of ozone, and not ozone itself, may be the oxidizing species, accelerating or decreasing the reaction rate. The idea that the hydroxyl radical may catalyze the decomposition of ozone is supported by the fact that ozone is reported to be more stable at high hydroxide ion where the hy(1) P. S. Bailey, "Ozonein Water and Waste Water Treatment",F. L. Evans, Ed., Ann Arbor Science, Ann Arbor, MI, 1972, pp 29-52. (2) B. D. Goldstein, C. Lodi, C. Collinson, and 0. J. Balchum, Arch. Enuiron. Health, 18, 631-5 (1969). (3) H. E. Stokinger, "Ozone Chemistry and Technology",American Chemical Society, Washington, DC, 1959, pp 360-9. (4) J. Hoigne, Proc. Int. Symp. Use High-Leuel Radiat. Waste Treat.-Status Prospects, 1975,410, 297-305 (1975). (5) J. Weiss, Trans. Faraday SOC.,31, 681 (1935). (6) H. Taube and W. C. Brav. J. Am. Chem. SOC..62. 3357 11940). (7) J. Hoign6 and H. Bader, Water Res., 10,377 (1976);'Vom Was&, 48, 283 (1977); Enuiron. Sci. Technol., 12, 79 (1978). (8)J. Staehelin and J. Hoign6, "Proceedingsof the 5th Ozone World Congress", International Ozone Association, West Berlin, April 1981. (9) J. Staehelin and J. Hoign6, private communication. (10) V. Rotbmund and A. Burgstaller, Monatsh. Chem., 34,665 (1913). (11) K. Sennewald, 2.Phys. Chem., Abt. A, 164, 305 (1933). (12) M. G. Alder and G. R. Hill, J. Am. Chem. SOC.,72,1884 (1950). (13) W. Stumm, H e h . Chim. Acta, 37, 773 (1954). (14) M. L. Kilpatrick, C. C. Herrick, and M. Kilpatrick,J . Am. Chem. Soc., 78, 1784 (1956). (15) G. Czapki, A. Samuni,and R. Yellin, Isr. J. Chem., 6,969 (1968). (16) C. G. Hewes and R. R. Davidson, MChE J., 17, 141 (1971). (17) L. Rizzuti, V. Augugliaro, and G. Marrucci, Chem. Eng. Sci., 31, 887 (1976). (18) S. Morooka, K. Ikemizu, and Y. Kato, Int. Chem. Eng., 19,650 (1979). (19) L. J. Heidt and V. R. Landi, J. Chem. Phys., 41, 176 (1964).
0 1982 American Chemical Society
256
The Journal of Physical Chemistry, Vol. 86,No. 2, 1982
Forni et al.
droxyl radical dissociates.22 Experimental Section Materials. Ozone was generated by an electric discharge in oxygen in an ozonator and adsorbed on silica gel at -78 "C. Stock ozone solutions which contained acetic acid (4 x mol dm-3) as a stabilizer were prepared as described.23 The carbonate-free alkaline solutions were obtained from a stock sodium hydroxide solution which was prepared by reacting sodium metal with water under a helium atmosphere." All solutions were prepared from degassed, deionized, Millipore-filtered water (RH0 I18 MO cm), the quality of which corresponds to triply distilled water. Sodium carbonate and sodium acetate were of analytical grade (Merck) and used without further purification. Procedure. Kinetic measurements and spectral identification of the intermediates were performed by using a Durrum Rapid Kinetics System D-110 stopped-flow spectrophotometer. The cell lengths employed were 2,10, or 20 mm depending on the conditions. The dead time of the system ranged from 4 ms in the 2-cm cell to approximately 1ms in the 2-mm cell. For the optical measurements, a 75-W xenon lamp with long-time stability of the base line was used. Two dispersing elements were employed: a prism in the 180-330-nm range and a grating from 330 to 900 nm. The wavelength setting of the monochromator was adjusted by using the well-defined spectrum of a Hg arc lamp. With pure water as a reference, the 0% and 100% transmission values were adjusted for each individual experiment (U, = 10 V f 0.1%). The time-constant setting of the instrument for traces obtained was kept, where possible, to below 10% of the reaction time constant keeping degradation of the trace rise time to a minimum. All results were recorded as transmission vs. time curves on a Tektronix 7633 storage oscilloscope and photographed. These data were transfered to a PDP 11/40 computer for kinetic analysis and processed as described elsewhere.25 Ozone and alkali solutions were introduced into the stopped-flow spectrophotometer via separate gas-tight syringes, acting as reservoirs. These syringes were linked to the two spectrophotometer drive syringes through Teflon tubing attached to the syringe valve block. The ozone solutions containing acetic acid were protected from heat and UV light, and no appreciable decay of the ozone was noted throughout an experiment. All experiments were performed at 20 f 1 "C. Results and Discussion Effect of [OH-]. Aqueous ozone solutions exhibit an ultraviolet absorption band with a maximum at 260 nm where e = 3314 70 dm3 mol-' cm-1.23 When carbonatefree ozone and alkali solutions are mixed, the decay of ozone can thus conveniently be monitored at this wavelength. The insert of Figure 1 shows, for example, the change in transmission a t 260 nm after mixing a solution mol dm-3 O3 with 0.01 mol dm-3 OH-. The of 3 X ozone decay under these conditions is a pseudo-first-order reaction with a half-life decreasing with increasing pH, as is illustrated in Figure 1. From the slope of the plot t l j 2 vs. [OH-I-l, a second-order rate constant for the overall
*
(20)L. J. Heidt, J. Chem. Educ., 43, 623 (1966). (21) V. R. Landi and L. J. Heidt, J . Phys. Chem., 73, 2361 (1969). (22)J. L. Weeks and J. Rabani, J. Phys. Chem., 70, 2100 (1966). (23)E.J. Hart, unpublished results. (24)D. Bahnemann and E. J. Hart, J.Phys. Chem., preceding.~ paper in this issue. (25)N. Shinohara, J. Lilie, and M. G. Simic, Znorg. Chem., 16,2809 (1977). ~
- LOO
t
200 100
0 '
100
200 300 LOO [dm3 mol-']
Figure 1. Plot of reaction half-life ( t i , 2 ) vs. reciprocal OH- ion concentration (l/[OH-]) for the decay of ozone at 260 nm in carbonatefree solutions. 8 = 20 f 1 O C . The apparent secondorder rate constant derived from the slope: k,,, = 540 dm3 mol-' s-'. Insert: Oscilloscope trace (transmittancevs. bme) obtained at 260 nm on rapid mixing of 3 X mol dm-3 O3with 5 X lo3 mol dm-3 OH-, both solutions carbonate free, 0 = 20 f 1 OC, U o = 10 V. The base line presents the transmittance of pure water. The reaction half-life is approximately 240 ms; cell length = 20 mm. L' I,
d2t 0-
1
I
i
I
1.
u
'350
400 450 A Inml
500
Flgure 2. Absorption spectrum observed after rapid mixing of 2.5 X I O 4 mol dm3 ozone with 0.05 mol dm3 OH- ions and 5 X lo-' mol dm-3 oxygen. 0 = 20 f 1 "C;cell length = 10 mm. Insert: Oscilloscope trace (transmittance vs. time) obtained at 430 nm on rapid mixing of 2.5 X lo-' mol dm-3 ozone with 0.05 mol dm-3 OH- and 5 X lo4 mol dm-3 oxygen. 8 = 20 f 1 OC;U o = 10 V. The base line presents the transmlttance of pure water; cell length = 20 mm.
*
decay of ozone is calculated as 540 20 dm3mol-' s-l. This value is lower than those of Czapski et al.15 and Rizzuti et al.," who reported biomolecular rate constants of 700 and 1500 dm3 mol-' s-', respectively. As O3 decays, the ozonide radical ion, 03--forms simultaneously. It has been identified by its spectral and kinetic properties. Favorable experimental conditions for the appearance and stabilization of the ozonide radical anion were chosen, i.e., high pH and added oxygen (see discussion below). A typical transmittance vs. time trace of the transient detected under these conditions (2.5 x lo4 mol dm-3 0,; 5 x mol dm-3 02,0.05 mol dm-3 OH-) at 430 nm is shown in the insert of Figure 2. The spectrum taken at maximum height of this absorption band (Figure 2) agrees satisfactorily with the established spectrum%of the ozonide ion in water and therefore identifies this transient species as 03--. Assuming that the ozonide radical anion is an immediate product of the reaction of OH- with ozone, a value of kl = 160 f 20 dm3 mol-' s-l is calculated from the initial buildup of 03--.This rate constant is considerably lower than the value obtained from the overall decay of 03. This indicates that 03-(26)G.Czapski and L. M. Dorfman, J. Phys. Chem., 68,1169(1964).
Hydroxide Ion Initiated Decomposition of Ozone
- 300 -E. 200 v)
4
100
7
y
The Journal of Physical Chemisrry, Vol. 86, No. 2, 1982 257
500
N 1
cco;- [mol dm-31 x103
Flgure 3. Reaction half-life for the complete decay of ozone at 260 nm vs. concentration of carbonate added to the hydroxide ion solution. C = 2 X 10"' mol dm-3;,C = 2.5 X lo-* mol dm3; 8 = 20 f 1 Insert: Oscilloscope trace (transmittancevs. time) obtained at 260 nm on rapid mixing of 2 X 10"' mol dm3 O3 with 2.5 X 8 = 20 f 1 OC; U o mol dm-3 OH- plus 5 X mol dm-3 C O:-. = 10 V. The base line presents the transmittance of pure water; cell length = 20 mm.
%.
10 I
loo0
u \*
io-6
Flgure 5. Apparent first-order rate constant, k (s-I), for the complete decay of ozone at 260 nm vs. concentration (logarithmic scale) of acetate (A)and carbonate (0)added to the OH- ion solutions. C o = 0.7 X 104-1.5 X lo-' mol dm3; C , = 2.5 X lo-* mol dm-3; = 20 f 1 O C .
8
absorption spectrum shown in Figure 4 is very similar to that reported for co3-..22i27 Since C O P is most likely formed via reaction 2, further support is provided for the OH. COS2- OH- C O P (2)
fi-.
+
500
600
io-L 10.~ c [ m o l dm-31
700
A[nml Flgure 4. Absorptlon spectrum observed after rapid mixing of 2.5 X 10"' mol dm-3 ozone with 3.2 X mol dm-3 OH- ions and 2.5 X mol dm3 carbonate. 8 = 20 f 1 OC;pH 10.5. Insert: Oscilloscope trace (transmittance vs. time) obtained at 600 nm on rapid mixing of 2.5 X 10"' mol dm3 ozone with 5 X IO3 mol dm3 O K and 2.5 X mol dm3 Cot-. 8 = 20 f 1 O C ; U o = 10 V. The base line presents the transmittance of pure water; cell length = 20 mm.
and/or its decay products participate in the chain decomposition of O3 in alkaline solution. Effect of Carbonate. The effect of the carbonate ion on the decay of ozone measured at 260 nm is illustrated in mol dm-3) is Figure 3. When an ozone solution (2 X mixed with alkaline solutions containing various concentrations of Na2C03,the half-life of the O3 decay increases with increasing carbonate concentration. The effect is greatest at relatively low concentrations of CO:-, whereas above 0.0015 mol dm-3 further addition of carbonate has relatively little effect. The oscilloscope trace (see insert of Figure 3) exhibits an induction period that is followed by a much faster decay. This genuine initiation period is absent in carbonate-free solutions (e.g., insert of Figure 1). The study of the carbonate effect was limited to CO2concentrations below 0.01 mol dm-3 because of distorted optical effects introduced by concentration gradients for solutions of higher concentration. The carbonate radical ion, C03-., forms on the rapid mixing of an O3 solution with an alkaline carbonate solution. Figure 4 displays the absorption spectrum of the species present at a steady-state concentration level (shown as a minimum in the insert of Figure 4). The corresponding kinetic trace taken at 600 nm exhibits the formation, the development of a steady-state concentration, and the decay of the transient COP radical ion. The
-
+
transitory appearance of OH. radicals during the OH-catalyzed decay of 03.This conclusion was earlier reached by Hoign6 and Bader,7B who showed that the relative rates with which organic substrates compete for reaction intermediates agree with their relative OH. radical rate constants. The decay of the carbonate radical ion does not follow pure second-order kinetics as would be expected from reaction 3 with k3 = 1.25 X lo7dm3mol-' sd. This suggests COS-. + COB-. products (3)
-
that the carbonate radical anion itself might react with ozone. Furthermore, the decay kinetics are complicated by the fact that under our experimental conditions ozone has completely decayed after about 1200 ms, whereas the carbonate radical anion has a lifetime of approximately 1500 ms. Effect of Acetate. Acetate also retards the initial decay of ozone. As can be seen from Figure 5, the measured decrease in rate is, however, smaller than in the presence of carbonate at the same OH- concentration (pH 12.3). Since acetic acid has been used as a stabilizer of ozone throughout the experiments even though the concentrations have been relatively small (normally in the range of 2 X 10+-10 X lo4 mol dm-3), the acetate ion present at high pHs still influences the rate of O3 decomposition. Our conclusion that the OH. radical is a transitory intermediate is also supported by these acetate studies. The carbonate ion is found to be about 3 times more effective as an inhibitor of the O3 decay than the acetate ion, which agrees fairly satisfactorily with the ratio of 4 for their relative OH. radical rate constants.28 Effect of Hydrogen Peroxide. Preliminary studies have been carried out on the reaction of ozone with H02-. Hydrogen peroxide greatly accelerates the decomposition of O3in alkaline solutions. Figure 6 shows the oscilloscope traces obtained on rapid miying of 1 x mol dm-3 O3 with alkali in the presence and the absence of H202 For reaction mixtures with a final pH of 10.9, the half-life of the O3 decay is 1.9 s in the absence of H202and 0.047 s with 5 X mol dm-3 H202. This effect is greatly en(27) G. E. Adams, J. W. Boag, and B. D. Michael, Proc. R. SOC. London, Ser. A, 289, 321 (1966). (28) R. L.Wilson,C. L. Greenstock, G . E. Adams, R. Wageman, and L.M.Dorfman, Znt. J. Radiat. Phys. Chem., 3, 211 (1971).
258
The Journal of Physical Chemistry, Voi. 86, No. 2, 1982
Forni et al.
0.20
1 1.1
no H202 0.15
2 0.10
\ \
ti
B
0.05
5 x 105mol dm" H,O, 500
-0
* L O O ms
Fbwr 8. (A) Oscilloscope trace obtained at 260 nm on rapid mixing of 1 X lo4 mol dm3 ozone with 1.6 X lo3 mol dm3 OH-. Reactlon half-time t,12 = 1.9 s; 9 = 20 f 1 O C ; cell length = 20 mm. (B) Oscllogaph trace obtained at 260 nm on rapid mixing of 1 X lo4 mol dm3 ozone with 1.6 X mol dm-3 OH- and 5 X hydrogen peroxide. Reaction hab-time t = 47 ms; d cell length = 20 mm.
20
mol dm3
= 20 f 1 O C ;
230
270
250
A[nml Flgm 8. Absorption s p n m observed after r a p mixing of solutions of 5 X mol dm ozone with 2.5 X 10- mol dm-3 hydrogen peroxide and 5 X lo-* mol dm3 OH-. d = 20 f 1 O C . Insert: Oscilloscope trace obtalned at 240 nm on rapid mlxlng of 5 X lo-' mol dm-3 ozone with 2.5 X lo-' mol dm3 hydrogen peroxide and 5 X mol dm-3 OH-. d = 20 f 1 O C ; cell length = 20 mm.
been provided by the intermediate appearance of COP in 03-C02-solutions, reaction 6 would appear to be established as the primary reaction. However, these intermediates may alternatively be produced by primary reaction 7. Because reaction 4 between HOB-and O3 is found to
1
O3 + OH-
'*\
0 10.~
-.*-'
_.
10.~
10.'
cH202[mo1dm-ll
as Flgure 7. Reaction half-life of the ozone decay at 260 nm, t a function of the hydrogen peroxide concentration. C o = 1 X lo4 mol dm-3; C , = 2 X lo-* mol dm3; 8 = 20 f 1
+
-
HOy H++ 0 2 - * (5) Rate Constant of the Initiation Reaction. The above results confirm the earlier conclusions7+' that ozone decay in alkaline aqueous solution is not a single-step process but may involve a chain reaction. Our results also confirm that carbonate and acetate ions are chain inhibitors. Since the OH. radical has been proposed as a possible intermediate5p7-9in the decay of 03,the following chain initiation step can be envisaged: O3 + OH03-.+ OH(6) Since 03-.has been observed directly and additional support for the transitory formation of the OH. radical has
-
(29)B. H.J. Bielski and J. M. Gebicki, 'Advances in Radiation Chemistry", Vol. 2, M. Burton and J. L. Magee, W.,Wiley-Interscience, New York, 1970,p 177. (30)B. H.J. Biebki, Photochem. Photobiol., 28, 645 (1978),and references cited therein.
HOz-
+ O2
(7)
be very rapidly producing 03-.(one of our observed intermediates), reaction 7 must also be considered. 03-produces OH- radicals on dissociation by31-34
03-' + 0-. + o2
'e.
hanced at higher pHs and higher HzOz concentrations. Figure 7 presents the effect of HzOzconcentration on the half-life of the decay of 1 X mol dm-3 O3 at pH 12.3. A long-lived absorbing species formed during the decay of alkaline 03-HzOz complicates the kinetics as observed at 260 nm. The insert of Figure 8 illustrates the formation and the decay of this long-lived species at 240 nm and pH 12.7. (5 X lod mol dm-3O3and 2.5 X lo4 mol dm-3 H20z). This relatively long-lived species has an absorption spectrum which is very similar to the published spectrum of the superoxide anion radical Oz-..798 Therefore, it is apparent that a rapid reaction takes place between the anion of H202,HOz-, and O3 forming 02-., possibly via the following reactions:m*m H02- O3 H02. + 03-(4)
-
0-. + HzO + OH.
+ OH-
(8)
(9)
Therefore reactions 6 and 7 are equivalent initiation reactions, and a means must be sought for differentiating between them. Firstly, we consider the consequences of initiation reaction 6 relative to COS-. formation. Two OH*radicals result via reactions 6, 8, and 9. Accordingly, reaction 2 yields two carbonate radical anions per O3 molecule consumed in reaction 6. If, however, reaction 7 is the initiation process, three O3 molecules react to produce two OH. radicals resulting in two COS-. ions in the 03-alkaline carbonate system. In this case an 02--is converted into 03-.via reaction 10. This reaction has been shown to
-
Oz-. + O3
03-.+ Oz
(10)
proceed with a rate constant of about 1 X lo9 dm3 mol-l s-1a35
Summarizing these results, two COP per O3 molecule consumed are produced through the initiation reaction 6 and only two COS-. per three O3 consumed in initiation reaction 7. Consequently quantitative measurements of C 0 3 ; formation and O3 consumption should be helpful in resolving the question of the initiation step. The initial rate of the ozone decay at 260 nm was determined as reported above in 0.008 mol dm-3 carbonate where an induction period was found (insert of Figure 3). While the rest of the ozone decay curve followed fmt-order kinetics, the initial rate was much slower (Figure 9). This induction period is presumed to be the region where the (31)D.Behar and G. Czapski, Isr. J. Chem., 6,43 (1968). (32)J. Rabani and M. S. Matheson, J. Phys. Chem., 70, 761 (1966). (33)G. V. Buxton, Tram. Faraday SOC.,65, 2150 (1969). (34)B.L.Gall and L. M. D o h , J. Am. Chem. Soc., 91,2199(1969). (35)K. Sehested, J. Holcman, and E. J. Hart, unpublished results.
The Journal of phvsical Chemistry, Vol. 86, No. 2, 1982 259
Hydroxide Ion Initiated Decomposition of Ozone
0.2 0.4 t [SI
0
Flgure 9. Absorption vs. time curve obtained for the decay of ozone at 260 nm on rapid mixing of a solution containing ozone (C = 9 X loJ mol dm4) and an alkaline solution containing carbonate?C, = 7.9 X lo3 mol dm3, C-2- = 5 X lo3 mol dms). 9 = 20 f 1 O C . The slope of the initial decay (constant of reaction 7.
- -1
was used to determine the rate
TABLE I : Apparent Rate Constants of the Hydroxide Ion-Ozone Reaction in Aaueous Solutionsa reactant concn io6x
series
(k0.15)
[O,l, mol dm-3
A
11.2 11.9 12.3 11.3 11.7 12.3
90 90 90 200 250 250
vH
103 x [C03*-Ir
5 5 5 4 4 4
mol dm'3
h . . dm3 mol-, s" ~
B
123 + 94 r 127 t 101 ?: 125 r 118 +
43 40 43 17 21 23
Values for h , , obtained at various hydroxide ion concentrations for the two methods described. Series A shows the results from the initial decay of ozine at 260 nm in systems containing a high concentration of carbonate. Series B shows the results from the initial buildup of the carbonate radical anion at 600 nm. See text for calculations of rate constant for reaction 7.
OH. radicals formed from reaction 6 and 9 are scavenged by C032-according to reaction 2. Thus, the initial rate of the ozone decay under these conditions provides data for calculating kl from which k, or k7 may now be derived after the question of which initiation reaction is resolved. This primary rate constant should be the same at different OHconcentrations, and thus various pHs were studied (Table I, series A). The results yield a value for kl of 109 f 43 dm3 mol-' s-', which is independent of pH in the range 11.2-12.3. Since it is difficult to define the exact length of the initiation period, a relatively large experimental error exists. Furthermore, the reaction between the COP radical and O3 might change the initial kinetics slightly. The value of kl was also calculated from the initial rate of formation of the carbonate radical anion at 600 nm, a rate which depends only on the rate of reaction 6 or 7, since
these are the rate-limiting steps in the reactions producing OH. radicals and subsequently C03-.. These results are also shown in Table I (series B) for different OH- concentrations. A mean value of kl = 113 f 20 dm3 mol-' s-' was determined by this method. The steady-state concentration of the carbonate radical anion reaches only 10% of the O3 concentration at its maximum. Further evidence regarding the initial step in the chain decomposition of ozone comes from estimates of the initial concentration of the carbonate radical anion formed. The yield of carbonate radical anion prior to the formation of the steady state (see insert of Figure 4) correlates with the concentration of ozone which decayed after the same time interval under identical conditions. For example, results at pH 10.5 using 2 X mol dm-3 ozone (reactant concentration) and 2.5 X mol dm-3 carbonate (reactant concentration) show that after 50 ms approximately 9 X lo4 mol dm-3 of ozone has decayed; thus, if reaction 6 is occurring, one would expect 18 X lo* mol dm-3 of OH. radicals or C03--since two OH. radicals are produced per ozone reacted. The yield of carbonate radical anion after the same time interval is approximately 7 X 10" mol dm-3. Thus, the ratio sf OH. and C03-. formation to O3 consumption is 0.78 in these experiments. This result supports reaction 7 as the initiation step instead of reaction 6. Further evidence that reaction 7 is the primary step is provided by the extensive studies of Hoign6 and Bader, who report that 0.55 f 0.08 mol dm-3 of OH. may be produced from 1 mol dm-3 of O3 at pH 10.5.7 On the assumption that reaction 7 is the initiation step, k7 may be calculated from the values of kl listed in Table I. Since three molecules of O3 are decomposed per initial reaction in series A, k7 = (114 f 42)/3 dm3 mo1-ls-l = 38 f 14 dm3 mol-' s-'. In series B, two molecules of COS-. form per initial reaction; therefore, k7 = (115 f 20)/2 dm3 mo1-ls-l = 57.5 f 10 dm3 mol-' s-'. From the buildup of the absorption of the 03-radical anion at 430 nm, a rate constant, k = 160 f 20 dm3mol-' s-', was calculated. Since two 03-.result from each initial step, k7 = (160 f 20)/2 dm3mol-' s-l = 80 f 10 dm3 mol-' s-l. This latter value, however, does not represent the true value of k7 since the ozonide ion is probably formed via reaction 10. Thus, the rate constant from the ozonide formation represents an upper limit for k7. For this reason we consider an average of the two values derived from the data of Table I (k7 = 48 f 12 dm3 mol-' s-l) as a more reliable rate constant of reaction 7. A value of 60 dm3 mol-' s-l has recently been reporteds for k7.
Acknowledgment. We thank Professor K.-D. Asmus for encouragement in undertaking this project and for discussions of the results. E.H. is also indebted to Dr. J. Hoigne and J. Staehelin of EAWAG, Dubendorf, Switzerland, for recommending carbonate and acetate as OH scavengers in the initial stages of this study and to K. Sehested and J. Holcman of the Riso National Laboratory, Denmark, for valuable discussions on the mechanism of the primary reaction. This work was performed during the third industrial training period of L.F. as part of his Bachelor's degrees on Applied Biochemistry, Brunel University, U.K. E.H.thanks the Alexander von Humboldt-Stiftung for a Senior U.S. Scientist Award.
