Environ. Sci. Technol. 1991, 25, 1589-1596
(8) Somasundaran, P.; Middleton, R.; Viswanathan, K. V. In StructurallPerformance Relationships in Surfactants; Rosen, M. J., Ed.; ACS Symposium Series 253; American
Chemical Society: Washington, DC, 1984; Chapter 17.
(9) Yeskie, M. A,; Hawell, J. H. J. Phys. Chem. 1988,92,2346. (10) Bisio, P. D.; Cartledge, J. G.;Keesom, W. H.; Radke, C. J. J . Colloid Interface Sci. 1980, 78, 225. (11) Lee, J. F.; Crum, J. R.; Boyd, S. A. Environ. Sci. Technol. 1989, 23, 1365. (12) Standard Methods for the Examination of Water and Wastewater, 16th ed; Greenberg, A. E., Trussel, R. R.,
Clesceri, L. S., Ed.; American Public Health Association (APHA): Washington, DC, 1985. (13) Vold, R. D.; Vold, M. J. Colloid and Interface Chemistry; Addison-WesleyPublishing Co. Inc.: Reading, MA, 1983.
(14) Anderson, P. R.; Benjamin, M. M. Environ. Sci. Technol.
1990, 24, 692.
(15) Chiou, C. T. Reactions and Movement of Organic Chemicals in Soils. SSSA Spec. h b l . 1989, No. 22. (16) Hassett, J. J.; Banwart, W. L.; Griffen, R. A. In Environment and Solid Wastes; Francis, C. W., Auerback, S. I., Eds.; Butterworth Woburn, MA, 1983; Chapter 15.
Received for review November 27, 1990. Revised manuscript received April 22,1991. Accepted May 14,1991. This work was supported in part through the Industrial Waste Elimination Research Center, a national research center sponsored in part by the U.S. Environmental Protection Agency Office of Exploratory Research.
The Primary Reaction in the Decomposition of Ozone in Acidic Aqueous Solutions Knud Sehested," Hanne Cotfltzen, Jerzy Holcman, Chrlstlan H. Flscher,t and Edwln J. Hart*
Environmental Science and Technology Department, Ris0 National Laboratory, DK 4000 Roskilde, Denmark The primary step in the thermal decomposition of acidic aqueous ozone solutions is the dissociation reaction: O3 F? 0 + 02.This reaction is supported by the retarding effect of O2 on the decomposition rate, by our activation energy measurements, E A = 79.5 f 8.0 k J mol-', and by isotopic exchange experiments with 18J802. The rates of ozone decomposition were measured at pHs 0-4, at temperatures 0-46 "C, and at O2concentrations 10-5-10-3 M. The 0 atom formed in the dissociation reaction is postulated to be the precursor for OH and H 0 2formation. The ionization of the H 0 2 free radical (pK = 4.8) accounts for the increase in rate with rising pH. The mechanistic implications of O3 decomposition in acid solution are briefly discussed.
Introduction The mechanism of aqueous ozone decomposition has been studied for over half a century (e.g., refs 1-23). Early work has shown that ozone decomposes faster in increasingly alkaline solution. The chain-reaction characteristics of the self-decomposition reaction of aqueous ozone have clearly been pointed out, and it is generally accepted that O3 decomposition is base-catalyzed. Previous work (10, 12,23,24)has established that in the presence of radical scavengers the initiation rate above pH 8-9 is proportional to the concentration of both O3 and OH-. One of the proposed initiation reactions in alkaline solutions is
O3 + OH-
-
H02- + O2
(1)
The rate constant of initiation reaction 1 is 40-70 dm3 mo1-ls-l (2,9,10,12), and it has been shown that OH, 02-, and 0, radicals are intermediates in the subsequent chain decomposition. Pulse radiolytic investigations (13-20)have recently elucidated new facets of the propagation and termination steps of the aqueous ozone decomposition by determining the rate constants of the free-radical reactions with ozone. In acid solutions, however, reaction 1cannot be the sole initiation reaction, since the observed reaction rates are +
t
Hahn-Meitner-Institut,D 1000 Berlin 39, Germany. Port Angeles, WA 98362.
0013-936X/91/0~25-1589$02.50/0
much higher than would be predicted from the low OHconcentrations below pH 4. Although an additional initiation reaction is necessary, definite identification of such a reaction has not been forthcoming. Our objective is to identify the primary decomposition reaction in acidic ozone solutions. An early proposal (23) postulated formation of OH radicals in a reaction with water: O3 + H 2 0 20H + O2 (2) This reaction has also been suggested in a more recent study (6). Experiments (25) with 180-enrichedwater in 0.04 M HCIOl have established little or no exchange between the oxygen atoms of O3 and water, but upon decomposition of O3 an exchange does take place. From the extent of exchange it was concluded that the active intermediate was the OH radical and not H02 or C104. In the present paper we report on the effect of O3 and O2 concentrations and temperature on the thermal decomposition of acidic ozone solutions in the pH range 0-4 and in the temperature range 0-46 "C. Tracer experiments with added 18Ja02 were also carried out in order to determine the importance of the 0 atom as a major initiating species.
-
Experimental Section Special care was taken to avoid the effects of impurities on the decomposition of ozone solutions. All glassware was baked at 450 "C for 4 h after cleaning and then flushed with the solution before use. The water used was triply distilled, and in most experiments preozonized. Other chemicals were reagent grade and were used without further purification. All experiments were carried out with reagent grade HC104. The pH was measured on a Radiometer PHM22s pH meter. The preparation of (2-4) X M O3 stock solutions, nearly free of 02,has been previously described (15-17). Stock solutions were also prepared from O3 gas stored in the dark in moist acidic 100-mL glass syringes at 0 "C. We found that O3 stored as gas under these conditions was much more stable than 0, dissolved in acidified aqueous solutions. However, in order to prevent explosionbf the O3 in the syringes, unusual handling precautions are required. We have found that an explosion of 50-75 mL of
@ 1991 American Chemical Society
Environ. Sci. Technol., Vol. 25, No. 9, 1991
1589
tI
Table I Determination of O2 Concentrations in 0, Solutions (lo-, M HCIO,) -, Usinn Acrvlic Acid for Removal of Residuai Ozone I
solution
A (1.0 M NaC104)
t, min
09,
0 146 60 120 1000
B (0 M NaC104)
02,
pM pM
97 81 37 0 144 60 86 120 66 IO00 16
28 89 132 215 37 116 156 246
-
pM
AOz,
pM
A02/A03
49 64 109
61 104 187
1.25 1.62 1.73
58 78 128
79 119 209
1.36 1.52 1.63
av
0
200
250
300 h nm Figure 1. Spectra of 255 MM O3In 0.01 M HCIO, (A) and of 255 pM M acrylic acid, 1-2 mln after mixing (B). 0,
+
>50% O3 will completely demolish a 100-mL syringe. Explosions may be initiated by contact of the gas with dry glass joints, by organic matter, and by static electricity (very dry laboratory conditions were particularly hazardous). We found it necessary to operate under the following conditions in order to avoid explosions: All O3 preparations and transfer of concentrated O3 gas were carried out in a safety hood with a heavy glass window. Standard practice operations for the handling of explosives were also followed. In addition, the floor of the hood was covered with aluminum foil, grounded by a stream of water. The foil in the collecting area was also covered with moistened paper. Moistened gloves, grounded before transfer of the O3 into experimental or analytical syringes, were worn. With these precautions, no explosions have occurred. No danger exists after the O3 has been dissolved in the aqueous solutions. Solutions containing various amounts of oxygen and ozone were obtained by diluting the stock solution with the appropriate amount of 02-Ar-saturated acid. The ozone concentration was determined spectrophotometrically by its optical absorption at 260 nm in solutions using our previously measured extinction coefficient of 3300 dm3 mol-l cm-' (14) rather than the earlier value of 2900 dm3 mol-' cm-' (26). After removal of O3 by acrylic acid (AA) (27),the O2remaining in the solutions was analyzed by gas chromatography (GC). With an excess of AA, the reaction is rapid and complete (note the complete absence of the 0, band at 260 nm in Figure 1). Then the dissolved O2 was extracted from the solution in the Van Slyke pipet and analyzed by gas chromatography. This method was tested in a series of decay experiments where O2 and O3 were measured as a function of the time in 1.0 mM HC104 solutions in the presence (A) and absence (B)of 1.0 M NaC104 (see Table I). At the indicated times, 20 mL of solution was withdrawn from a 100-mL syringe. Ten milliliters was used for the 0, spectroscopic analysis and 10 mL for the O2 gas chromatographic analysis (28). In the gas analysis procedure, 0.25 mL of 0.1 M AA was injected into the 10-mL syringe to remove the 03.Within the 5-10 min required for the extraction of this solution on the Van Slyke apparatus, no liberation of 0 2 due to reaction product of 0, and AA occurs. Although the individual ratios, A 0 2 / A 0 3 in Table I, show a certain systematic variation, the average is close to the expected ratio of 312. On the basis of these experiments, we assume that our method for removing O3 with acrylic acid is suitable for measuring O2 in the presence of O3 in aqueous solu1590 Environ. Scl. Technol., Vol. 25, No. 9, 1991
1.52
Table 11 Initial O2Concentrations in Oxygen-Deficient Ozone Solutions (pH 2.0) 03,
PM
59 220 285 324 441 489 496 510 519 549 569
av
02, PM
Odo3
4 12 22 11 46 61 58 28 28 40 51
0.061 0.056 0.076 0.034 0.103 0.125 0.104 0.054 0.054 0.072 0.083 0.075
tions. We calculate the O2concentration at time t , [021t, of all decayed solutions from the initial oxygen concentration, [02]i, and the change in ozone concentrations, A[ O,],by the equation = [Odi + 3/2A[o3I (3) For experiments where low contents of O2 were required, we prepared 50-75% pure O3 gas with O2as the principal impurity by eluting O3 adsorbed on cooled silica gel with N20or C02 (14). The gaseous O3was collected and stored in 100-mL syringes, moistened with 0.01 M HC104acid and cooled to 0 "C. The 02-O3 solutions were prepared by adding 10 mL of the O3 gas to 100 mL of 0.01 M HC104 solution, inverting 10 times, expelling the remaining O3gas, and analyzing the resulting solution for O2and 0, by the AA method described above. Typically, 0.25 mL of 0.1 M AA was added to the extracted 10-mL solution within 3 min of withdrawal from the 100-mL syringe. Some representative data for the initial 0, and O2 concentrations are given in Table 11. The ozone to oxygen ratio increases in the solution from that in the gas phase (4) because of a 15-fold higher solubility of ozone in water. Acidic O3 solutions were kept in 100-mL glass syringes sealed with ground-glass capillary tips. The syringes were covered with aluminum foil in order to prevent photolysis from the laboratory light, and they were heated in a water bath at various temperatures maintained at fl "C. The syringes were wrapped in a plastic bag to prevent direct contact with the water. To evaluate the importance of possible wall reactions, the O3 decomposition rates were measured in seven different reaction vessels with surface to volume ratios from 415 to 2175 cm2/L. No significant difference in O3 decay was observed. It is therefore concluded that wall reactions are unimportant for the rate measurements. The photolysis experiments were performed with a Philips 25-W low-pressure sodium lamp (A = 589.2 nm)
- ---
peaksof runl. peaksof run 2
4
28
I I mass
A
attenuation
I
I
2
1
3
4
x io4 s
Flgure 3. First-order kinetic plot of ozone decomposition; 200 M ozone in 0.01 M HCIO, at 31 "C. Key: 0, lo-^ M 0,; X, 10 M
-!
02.
injectionl. 0
0.5
1.0
1.5
D
rnin
injection2. I m 0 0.5 1.0 1.5 rnln Flgure 2. Mass chromatogram of experiment 2, Table IV; retention time for O2 1.4 mln. (See text for details.)
irradiating the solution at a distance of 10 cm in a Pyrex syringe. A control solution in a syringe wrapped in aluminum foil was simultaneously irradiated. In the tracer studies with '*J802 special care had to be taken in mixing O3 and 18J802 in order to avoid isotopic exchange in the gas phase. We found that 1sJsJ603 and ' ~ 1 8 0 2could not be mixed in the gas phase without isotopic exchange. Consequently, the ozone and oxygen solutions were separately prepared with degassed water and stored in 100-mL syringes. Mixing of the two solutions was performed by transferring the solutions from syringe to syringe through a capillary glass connector. The isotopic exchange during thermal decomposition was carried out in 100-mL syringes in the dark. At suitable intervals, 25 mL of solution was withdrawn from the heated syringes and quickly cooled to room temperature. Five milliliters was used for the spectrophotometric O3determination. Then 0.1-0.25 mL of 0.1 M acrylic acid was injected into the remaining 20 mL of the cooled solution in order to destroy the residual ozone. Next, this solution was drawn into the pipet of the Van Slyke manometric gas analyzer, and the gases were extracted by equilibration with a 30-mL gas phase. The gases were then withdrawn and stored in 3-mm-diameter Pyrex tubes, which were sealed, top and bottom, with short columns of mercury. These tubes contained of the order of 0.5 mL of gas. Gas chromatographic and mass spectrometric analyses were carried out on 0.05-0.1-mL samples of the extracted gas. Some analyses revealed from 10 to 30% air and N20 contamination in the samples. The fraction of these impurities was estimated and corrected for. N2and 0, were separated and analyzed on a molecular sieve column using an integrating Delsi gas chromatograph operating at 80 "C. N,O and C02 were separated on a Porapak column. In the mass spectrometer oxygen and nitrogen were fiist separated in a GC apparatus on a 5-A molecular sieve column at 175 "C and 63 psi with He as carrier gas (50-pL sample volume). The mass spectrometer was operated in the Multi-Ion-Selection mode (mass chromatogram), observing the masses 32,34,36, and 28 (nitrogen) as a control for air contamination (29) (see Figure 2). The actual peak height is hobs X attenuation. When a nitrogen peak was detected, a quarter of its height was subtracted from the
mass 32 peak. The ratio of the peak heights of masses 32(corrected for air):34:36 gave the isotopic ratio. At least two injections per sample were made. Figure 2 shows two runs on the same sample. The times of injections are marked. The second sample was injected before the first had eluted. Reproducibility was better than 5%. The negative peak on mass 28, located just above the oxygen peaks, is due to the fact that we have a certain background of nitrogen even in high vacuum. When another gas, here oxygen, enters the source, the partial pressure of nitrogen decreases. Water is almost completely adsorbed on the molecular sieve column. At the retention time of oxygen no water from the sample is present in the ion source and no fragments of it disturb the 0, result. Results Kinetic Treatment. Although the thermal decomposition of ozone in aqueous acidic solutions is a chain reaction and does not follow a simple kinetic order (22),we have used first-order kinetic plots to obtain the initial rates of the O3decomposition. The change in the initial rate, expressed as a first-order rate constant, is used as the measure of the effect of O3 and O2 concentrations, pH, temperature, and visible light on the O3decomposition. The initial rate is measured over approximately 1half-life of the O3 decomposition, where the deviation from firstorder kinetics is within the overall accuracy of our experiments (&lo%). As an example, the thermal decomposition at 31 "C of an oxygen-depleted (10 pM 0,) and an oxygen-saturated (lo3 pM 0,) 200 pM O3 solution in 0.01 M HC10, is shown in Figure 3. The initial first-order rate constants for the two solutions are kl = 1.3 X lo4 s-' and k2 = 4.7 X s-', respectively. The reproducibility was better than 5%. 0,Concentration Effect. The initial decomposition rate depends on the ozone concentration. The first-order rate constant increases by a factor of -1.6 when the O3 concentration is doubled. Figure 4 shows the effect of O3 concentration on the first-order decay plots. These runs were carried out at 31 "C, pH 2.0, on 30-2000 pM O3 solutions containing -500 pM 0, initially, and the solid curves are computed from a tentatively proposed mechanism (see Discussion). Similar decay curves and 0, concentration dependencies are found at all pHs between 0 and 4. Oxygen forms during the course of ozone decay according to the overall reaction
-
203 30, (4) Millimolar O3solutions become supersaturated with oxygen during decay of the 03. We have measured supersaEnviron. Sci. Technol., Vol. 25, No. 9, 1991 1591
Table I11 Initial First-Order Rate Constant as a Function of O3and O2 Concentrationsn
09 08 0.7
06 0.5 01 0
"TO3
0.2
c 01
h 2 '
'
'
I
2
1
I
\ \
,
,
3
,
j
5
L
x io4 s Flgure 4. Relative firstorder decay (C/Co)of 0,solutions containing 5 X lo4 M O,, pH 2,at 31 O C . Solid lines are computed according to the proposed mechanism (see Discussion). Key: 0 , 30 pM ozone; 0,60 pM ozone; X, 124 pM ozone; A g220 pM ozone; 0,450 pM ozone; +, 970 pM ozone; V, 2070 pM ozone.
03, PM
02, &f
temp, " C
los ki, s-l
93 88 91 95 199 209 200 202 361 382 349 379 146 285 655 146 286 669
8 104 402 1020 10 105 390 1020 25 120 420 1020 7 15 36 1150 1150 1150
31 31 31 31 31 31 31 31 31 31 31 31 25 25 25 25 25 25
8.6 6.2 4.8 2.4 13.2 9.6 7.0 4.7 18.2 13.5 11.5 8.0 4.0 6.3 10.5 1.3 2.6 5.3
"Data for Figure 6.
I
0
A
1
I
t
L
200 LOO 600 800
500
1000
PM02 Flgure 6. Effect of oxygen concentration on the initial firstorder rate constant of 0,decomposition at pH 2 (see Table 111). (A) The ratio of the rate constants in 0,depleted and 0,-saturated (lo-, M 0,) solutions as a function of ozone concentration. Key: 0, 31 OC, X, 25 O C . (B) Initial first-order rate constant for three different O3concentrations at 31 OC as a function of the O2concentration (see Table 111). Key: X, 90 pM Os;0,200 pM 0,;0,370 pM 0,. N O 3
0
1
2 PH
3
4
Flgure 5. Effect of pH on the initial firstorder rate constants for 0, solutions at 0 and 31 O C containing various amounts of 0,.Key: 0, 285 pM O,, 5 X 104M 0,'31 O C ; 0,33 pM O,, 5 X lo4 M O,, 31 O C ; X, 300 pM 0 ,,2.5 X lo-' M O,, 0 O C ; A,225 pM O,, 1.4 X lo4 M O,, 0 OC.
turations of 2-3 mM oxygen. Gas bubbles develop in these solutions, and because of these bubbles some uncertainty in our ozone determinations at high O3 concentrations exists. However, calculations have shown that supersaturation has a negligible effect on the determination of the initial rate constant. Effect of pH. The ozone decay rate varies with pH in the acid range. The initial decomposition rate increases by a factor of -2 for each pH unit in the range 0-4. The effect of varying O3and O2concentrations in this pH range is given in Figure 5. Here the first-order initial decomposition rate constants for O3solutions at 31 O C containing 500 pM O2and for O3solutions with 25 and 1400 pM O2 1592 Environ. Sci. Technol., Vol. 25, No. 9, 1991
at 0 O C are shown as a function of pH. The slopes of the In ki vs pH curves are identical. Since the O3concentration was also varied in these experiments, we conclude that the effect of pH on the initial decomposition rate is independent of both O3 and O2concentrations as well as temperature. Oxygen Concentration Effect. Oxygen stabilizes acidic aqueous O3solutions. Several examples are shown in parts A and B of Figure 6 (data in Table 111). In solutions fairly free of oxygen, e.g., solutions containing
