Environ. Sci. Technol. 2003, 37, 1941-1948
Photolysis of Ozone in Aqueous Solutions in the Presence of Tertiary Butanol ERIKA REISZ,† WINFRIED SCHMIDT,‡ HEINZ-PETER SCHUCHMANN, AND C L E M E N S V O N S O N N T A G * ,§ Max-Planck-Institut fu ¨ r Strahlenchemie, Stiftstrasse 34-36, P.O. Box 101365, 45470 Mu ¨ lheim an der Ruhr, Germany
The ozone decomposition quantum yield (Φ) in millimolar and higher-concentration aqueous tertiary butanol solution is 0.64 ( 0.05 (observed over a wavelength range from 250 to 280 nm) and rises toward lower tertiary butanol concentrations (Φ ≈ 1.5 at 10-5 M at pH 2) on account of the onset of the well-known •OH-radical-induced chain reaction. The destruction of the organic is initiated by hydrogen-atom abstraction through OH radicals which are produced via the reaction of the photolytically generated O(1D) with the solvent water at a quantum yield of Φ(•ΟΗ) of about 0.1. There is no decomposition of ozone in the dark on the time scale of the photolysis experiment. The efficiency of tertiary butanol destruction with respect to ozone consumption ([O3]o ) 3 × 10-4 M), defined by the ratio ∆[t-BuOH]/∆[O3], termed η(t-BuOH), is 0.26 at millimolar tertiary butanol concentrations, determined at the stage of essentially complete ozone consumption. It diminishes toward lower tertiary butanol concentrations (∆[t-BuOH]/∆[O3] ≈ 0.17 at [t-BuOH]o ) 1 × 10-4 M). Part of the effect of the ozone, apart from being a source of •OH radicals, rests on the intervention of HO2•/O2•- which is produced in the course of the peroxyl-radical chemistry of the tertiary butanol in this dioxygen-saturated environment and converted into further •OH radical by reaction with ozone. Moreover in this system, organic free radicals and peroxyl radicals react with the ozone. On the basis of the experimental and mechanistic-simulation data, the quantum yield of direct (by hv) ozone cleavage in aqueous solution is estimated at about 0.5.
Introduction Ozone is widely used as an oxidant in pollution abatement in aqueous media. It has the advantage that it is an environmentally friendly oxidant as its use is residue-free. However, with many pollutants it reacts so slowly [for a compilation of rate constants see ref 1] that for all practical purposes they must be considered ozone-refractory. For this reason, procedures are being applied where the oxidative * Corresponding author phone: +49 208 306 3694; fax: + 49 208 306 3951; e-mail:
[email protected]. † On leave from the University “Politehnica” of Timisoara, Faculty of Industrial Chemistry and Environmental Engineering, Bul. Victoriei 2, 1900 Timisoara, Romania. ‡ Present address: Fachhochschule Gelsenkirchen, Neidenburger Strasse 10, 45877 Gelsenkirchen, Germany. § Present address: Institut fu ¨ r Oberfla¨chenmodifizierung (IOM), Permoserstrasse 15, 04318 Leipzig, Germany. 10.1021/es0113100 CCC: $25.00 Published on Web 03/22/2003
2003 American Chemical Society
power of the ozone is “activated”, i.e., by a combination of ozone with H2O2 (reaction 1) (2-6) or UV radiation (reaction 2) (7-14) which gives rise to the formation of OH radicals and other reactive species.
2O3 + H2O2 (HO2-, H+) f 2•OH + 3O2
(1)
O3 + hν + H2O f •OH + other species
(2)
The use of ozone together with adjuvants such as UV radiation is usually classed among the procedures termed “Advanced Oxidation Processes” (AOP) (2, 15) which are in general based on the production of the very reactive •OH radical which induces free-radical-based oxidation of organic solutes. Its reactivity exceeds that of ozone itself [for • OH-radical-reaction rate constants see ref 16]. Upon photolysis, O3 is decomposed into O2 and oxygen atom O(1D) and O(3P) (17-21). Below 300 nm in the gas phase, the main process is the formation of O(1D) (quantum yield, Φ ≈ 0.9) and singlet dioxygen, O2(1∆g), as well as dioxygen in its ground state, O2(3Σg-) (reactions 3 and 4). In comparison, the competing reactions 5 and 6 which yield O(3P) are of smaller importance (Φ ≈ 0.1) (17, 21-23). Φ(O(1D)) falls to a value near 0.1 above a wavelength of about 320 nm but not to zero. This shows that the spin-forbidden formation of O(1D) and O2(3Σg-) (reaction 4) is possible (19-21, 24).
O3 + hν f O(1D) + O2(1∆g)
(3)
O3 + hν f O(1D) + O2(3Σg-)
(4)
O3 + hν f O(3P) + O2(3Σg-)
(5)
O3 + hν f O(3P) + O2(1∆g)
(6)
O(1D) is very energetic [heat of formation, 437 kJ mol-1, cf. ref 25] and therefore reacts fast (k in the order of 109-1010 M-1 s-1) with practically all conceivable substrates including water [reaction 7, k ) 1.8 × 1010 M-1 s-1 (26)]. It is believed that O(1D) reacts predominantly by insertion into the C-H [cf. ref 27] or the O-H bond (28). In the gas phase, the excess energy of the H2O2 molecule so formed results in the fragmentation of the O-O bond (BDE ) 210 kJ mol-1 (29), reaction 8).
O(1D) + H2O f H2O2 (hot)
(7)
H2O2 (hot) f 2•OH
(8)
In aqueous solution, these reactions are expected to occur as well, but in the condensed phase, the rapid thermalization of the “hot” H2O2 (reaction 9) interferes with its break-up into two OH radicals (reaction 8). Moreover insofar as fragmentation does occur, the solvent-cage effect inhibits the escape of •OH, i.e., there is recombination, in competition with diffusion into the bulk followed by reaction with a solute. Thus, most of the O(1D) is expected to be converted into H2O2.
H2O2 (hot) f H2O2
(9)
Most of the O(3P) disappears by reacting with the groundstate dioxygen O2(3Σg-) introduced into the solution together with the ozone at a close-to-diffusion-controlled rate (30) VOL. 37, NO. 9, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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[reaction 10, k ) 4 × 109 M-1 s-1 (31)]. Its reaction with ozone is relatively unimportant under the present conditions [reaction 11, k ) 5 × 106 M-1 s-1 (32)]. The reaction of O(3P) with water, in contrast to that of O(1D), is endothermic by about 75 kJ mol-1 and is therefore nonexistent under the present conditions (30) (k ≈ 10-3 M-1 s-1 (33); extrapolation of high-temperature gas-phase data in ref 34 to room temperature leads to a value in the order of 10-1 M-1 s-1).
O(3P) + O2(3Σg-) f O3
(10)
O(3P) + O3 f 2O2
(11)
Regarding the reaction of O(3P) with organic material, rate constants in the order of 106-108 M-1 s-1 are expected for H-atom abstraction by O(3P) from -CH3 groups (cf. reaction 12) (35, 36). Taking neopentane as a model compound for tertiary butanol, the rate constant would be 2.5 × 106 M-1 s-1 (36). Therefore, under conditions of oxygen saturation, O(3P) will contribute to the degradation of the organic solute RH only at near-molar organic-solute concentrations (usually not realized in situations of water purification), in which case, at the same time as it reacts with the organic solute, it serves as a source of further •OH. The rate constants of the reactions of •OH with organic solutes (e.g., reaction 13) are generally large; for RH ) tertiary butanol, k ) 6 × 108 M-1 s-1 (16).
O(3P) + RH f •OH + R• •
OH + RH f H2O + R•
(12) (13)
In his pioneering work, Taube has concluded that only H2O2 is formed upon the photolysis of O3 in water (28). This conclusion was based on the apparent 1:1 stoichiometry of ozone consumption and H2O2 formation in the presence of acetate as an OH radical scavenger. At this time it was not yet known that in its •OH-induced reactions in the presence of O2 acetate gives rise to relatively large amounts of H2O2 (37). This implies that the 1:1 stoichiometry is in fact coincidental. In the present study, the yield of tertiary butanol disappearance which characterizes the efficiency of ozone use, and of several oxidation products, has been determined. The yield of tertiary butanol disappearance is equal to the yield of 2-methyl-2-hydroxypropyl free radical formed initially from tertiary butanol by the action of •OH [plus, at near-molar tertiary butanol concentrations, of O(3P)], in competition with its reaction with ozone (reaction 14, see below). Other free radicals occurring in this system are not reactive enough to react with tertiary butanol. The efficiency of ozone use, ∆[t-BuOH]/∆[O3], depends on the ratio of the initial concentrations of the cosolute and ozone as well as generally on the propensity of the cosolute to liberate HO2•/O2•- in the course of its oxidation (see below). This aspect is a consideration in the valuation of the O3/UV process, compared with other AOP.
Experimental Section Tertiary butanol solutions were made up in Milli-Q-filtered (Millipore) water, phosphate-buffered (pH 2) or neutral. Before irradiation the solutions were bubbled with ozone from a dioxygen-fed ozone generator (Philoz 2, Philaqua, Gladbeck, Germany) for several minutes, such that the final ozone content was ∼3 × 10-4 M and this was determined spectrophotometrically, taking (260 nm) ) 3300 M-1 cm-1, cf. refs 38 and 39 (for the ozone UV spectrum see Figure 2). The O2 content of these solutions is close to saturation (∼1.2 × 10-3 M). Experiments were carried out at room temperature 1942
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FIGURE 1. Photolysis of ozone at 254 nm in aqueous solution (pH 2) in the absence (O) and presence (0.01 M, B) of tertiary butanol. The logarithm of the ratio of the ozone concentration (c) at a given photolysis time (intensity, 5 × 1017 quanta s-1 dm-3) relative to that at time zero (c0, 3 × 10-4 M) is plotted as a function of the irradiation time. The dotted line indicates the ozone thermal (dark) degradation rate, in the presence of 0.01 M tertiary butanol.
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 9, 2003
FIGURE 2. Quantum yields at various wavelengths, of the decay of ozone in aqueous solutions (pH 2) containing 1 × 10-2 M tertiary butanol. The initial ozone concentration was typically 3 ×10-4 M. The solid line shows the ozone absorption spectrum. (20-21 °C). In the presence of 1 × 10-2 M tertiary butanol in the dark, half of the ozone remained after 16 h. A series of experiments was also carried out with neutral solutions where the half-time in the dark was found to be not materially shorter (∼15 h). For UV irradiation at 254 nm, a low-pressure mercury lamp [Heraeus Sterisol NN 30/89, output at 254 nm ∼90% of UV/Vis (40)] was used which does not emit 185 nm radiation [this would photolyze tertiary butanol (41)]. Actinometry was done using the ferrioxalate system (42, 43). The contribution to the bleaching of the ferrioxalate by wavelengths longer than 254 nm was determined by a separate measurement, blocking the UV light with a platelet of ordinary glass. The effect of the 254 nm radiation was then obtained from the difference (5 × 1017 photons s-1 L-1). Irradiations were carried out in Suprasil cells (Hellma), 4 × 1 cm2 (face) × 1 cm, and lasted in the order of a few minutes. The decrease of the ozone concentration was determined by measuring the optical absorption at the maximum of the ozone absorption spectrum (260 nm). Ozone photolysis
quantum yields were calculated from low-conversion data (up to about 20%). For wavelengths other than 254 nm (247, 265, 270, 280, 290, and 300 nm), a homemade electronically integrating actinometer similar to that described in ref 44 was used. Its light source was a high-pressure mercury lamp; only those wavelengths were selected where the emitted intensity was high. This precluded experiments around 254 nm since this spectral line suffers self-absorption in the high-pressure mercury lamp. The slit width of the monochromator was 3 mm, leading to a bandwidth of about 10 nm. The reading of this setup at a given wavelength is proportional to the number of photons absorbed. For the determination of the coefficient of proportionality, ferrioxalate actinometry was done at the selected wavelengths. The light intensity of this lamp varies with the wavelength, from 2 × 1017 to 3.5 × 1018 photons s-1 L-1, i.e., 8.0 × 1014 to 1.4 × 1016 photons s-1 per cell volume over the wavelength range investigated. Formaldehyde was determined by HPLC as its 2,4dinitrophenylhydrazone (45) or spectrophotometrically using the Hantzsch reaction (46). Other aldehydic products such as 2-hydroxyisobutyraldehyde were analyzed by the former procedure (column Zorbax X∆B-C18, diameter 4.5 mm, length 15 cm, eluent acetonitrile/water 50/50 v/v, flow rate 1 mL min-1). Hydrogen peroxide was determined with molybdate-activated iodide (47). Irradiations were also carried out at neutral conditions which allowed the determination, by gas chromatography (column Stabilwax 30 m), of the consumption of tertiary butanol. Acids (formic, 2-hydroxybutyric, and acetic) were analyzed by ion chromatography (Dionex DX-100, column AS 9, eluent 4.5 × 10-4 M Na2CO3). The aldehydes and acids were identified by comparison with authentic material. The products were determined from the solutions when essentially all the ozone had reacted. Any remaining traces of ozone were stripped out with oxygen before analysis. The production of acid in initially neutral solutions was also indicated by a drop of the pH. Acetone is formed as well but its quantitative determination is disturbed by the uptake of this compound into the solutions from the ambient laboratory air and has therefore not been done. Kinetic modeling was done using a program (48) based on the Gear integration method.
Results and Discussion The photolysis of ozone in aqueous solution has been studied at various wavelengths in the range of 245-300 nm, in the absence and presence of tertiary butanol. As can be seen from Figure 1, the decrease of the ozone is slower in the presence of tertiary butanol than in its absence where, however, the presence of water-borne impurities is likely. Under these conditions, the thermal (dark) reaction of ozone (dotted line in Figure 1) during the time required for the photolysis experiments is negligible in comparison (t1/2 in the order of 10 h) and thus much slower than in the absence of an organic cosolute, cf. ref 49, unless this can react directly with ozone. Similar data have been obtained at the other wavelengths. In the absence of tertiary butanol and any other OH-radicalscavenging material (conditions not explored in this work where in the absence of tertiary butanol, a quantum yield of ozone destruction of ∼3 was observed at a pH of 2), the quantum yields would be far above unity, owing to the OHradical-induced ozone decomposition chain reaction (50, 51). At a tertiary butanol concentration of 10-5 M, the quantum yield of ozone decomposition approaches a value of 1.5 (cf. Figure 3), i.e., under these conditions the ozone is destroyed in a short reaction chain consisting of reactions 14-17 and 19-22, which is initiated by •OH radicals formed in the primary process (cf. reaction 7 followed by 8). Some ozone is restored in reaction 10, given that under the
FIGURE 3. Quantum yield of ozone decay in the photolysis at 254 nm (initial ozone concentration, 3 × 10-4 M, pH 2) as a function of the tertiary butanol concentration down to 10-6 M tertiary butanol. present experimental conditions the O2 concentration is ∼1.2 × 10-3 M.
O3 + •OH f HO2• + O2 •
•-
HO2 a O2 •-
O3 + O2
+H •-
f O3
•
(14)
+
(15)
+ O2
(16)
•
O3 + HO2 f OH + 2O2
(17)
The rate constant for reaction 14 has been reported at 1.1 × 108 M-1 s-1 (16). The chain is propagated by O2•- [reaction 16, k ) 1.6 × 109 M-1 s-1 (1)]. When the experiments are carried out at pH 2 where the less-reactive HO2• radical predominates [equilibrium 15, pKa(HO2•) ) 4.8 (52)], the O2•reaction is inhibited. The rate constant of the reaction of HO2• with ozone in the gas phase (reaction 17), at 1.2 × 106 M-1 s-1 (32), is much smaller. In any case in aqueous solution at pH 2, the effective rate constant solely due to the presence of O2•- in the equilibrium is also in the order of 106 M-1 s-1. A contribution to the decomposition through induction by OH- [cf. reaction 18, k ) 48 M-1 s-1, or in the order of 100 M-1 s-1 (1, 53)] is precluded under these conditions.
O3 + OH- f O2 + HO2-
(18)
The O3•- radical is very unstable with respect to O2 elimination (54) [reaction 19, k ) 2 × 103 s-1 (55)]. The formation of O•- is followed by rapid protonation [reaction 20, cf. pKa(•OH) ) 11.8 (16)]. Moreover at the conditions of these experiments, the effect of the protonation of O3•- must be taken into account (reaction 21). The HO-OO• bond energy has been estimated at a value as low as ∼4 kJ mol-1 (56), i.e., at room temperature the lifetime of HO3• may be as short as to be in the order of ns, to decay with the liberation of •OH (reaction 22).
O3•- f O•- + O2 •-
O
•
+ H2O f OH + OH
(19) -
(20)
O3•- + H+ f HO3• •
(21)
•
HO3 f O2 + OH
(22)
Tertiary butanol which is practically unreactive toward ozone itself [k e 3 × 10-3 M-1 s-1 (1)], cf. Figure 1, acts as VOL. 37, NO. 9, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Degradation of Tertiary Butanol by O3/UV (254 nm) in Initially Neutral (pH ∼ 6.5) or Acidic (pH 2; Phosphate-Buffered 0.05 M) Aqueous Solutiona [BuOH]o
10-1 M
10-2 M
10-3 M
10-4 M
η(- BuOH) η(CH2O) η(CH2O), pH 2 η[HC(O)C(CH3)2OH] η[HC(O)C(CH3)2OH], pH 2 η(HCO2H) η(HCO2H), pH 2 η[HOC(CH3)2CO2H] η[HOC(CH3)2CO2H], pH 2 η(CH3CO2H) η(CH3CO2H), pH 2
nd 0.094 0.071 0.082 0.059 0.008 0.01 0.016 0.02 0.011 0.03
nd 0.083 0.054 0.086 0.053 0.005 0.007 0.015 0.008 0.008 0.02
0.26 0.07 0.057 0.08 0.051 0.006 nd 0.015 nd 0.009 nd
0.17 0.042 0.045 0.035 0.039 nd nd nd
Efficiency η(S) ) ∆[S]/∆[O3] of disappearance or formation of a species S after practically complete disappearance of the ozone (initial concentration, 3 × 10-4 M). Duration of the photolysis, ∼ 5 min. Values determined experimentally (averaged). nd. ) not determined. a
an •OH radical scavenger [reaction 23, k ) 6 × 108 M-1 s-1 (16)] and so inhibits this chain reaction. •
•
(CH3)3COH + OH f CH2C(CH3)2OH + H2O (23) Quantum yields obtained at different wavelengths in the presence of 0.01 M tertiary butanol are given in Figure 2. Beginning near 280 nm, there is a decrease in the quantum yield of ozone disappearance toward longer wavelengths. A similar behavior is observed in the gas phase at wavelengths near 300 nm where this phenomenon is ascribed to the crossing of the threshold, near 300 nm, to the outside of the spin-allowed reaction channel (cf. Introduction). The average plateau value of Φ(-O3) below ∼270 nm is 0.64 ((0.05). This is similar to a value of 0.62 (at 254 nm) in 0.06 M acetic acid/0.05 M perchloric acid solution (28). A lower value of 0.48 ( 0.06 has also been reported (57). The variation of the quantum yield as a function of the tertiary butanol concentration down to 10-6 M is shown in Figure 3. It is seen that toward low scavenger concentrations a chain reaction comes increasingly into play; at still lower tertiary butanol concentrations quantum yield values become less meaningful, depending on the concentration of OHradical-scavenging trace impurities originally contained in the water employed. With increasing tertiary butanol concentration the quantum yield drops, and at a tertiary butanol concentration above 10-4 M it reaches a plateau at Φ ≈ 0.64. The reaction system can be characterized by the efficiency η(S) ) ∆[S]/∆[O3] of the degradation or the formation of a species S after the disappearance of the ozone. This quantity depends on the initial concentrations of the ozone and its organic cosolute. Table 1 lists the η values of tertiary butanol disappearance and of some products. The ratio η(CH2O)/η(-tBuOH) reflects the “stoichiometry” of formaldehyde formation per •OH radical reacted. In the case of millimolar tertiary butanol this is 0.26. A very similar value of (30(5)% has been observed in the O3/H2O2/ tertiary butanol system under otherwise similar conditions (58). The chemistry of the system is determined mainly by the formation of peroxyl radicals, by their termination reactions, and by elimination reactions undergone by some of them, but also by the reaction of these and of carbon-centered radicals with the ozone. The free-radical chemistry of tertiary butanol induced by the •OH radical in the presence of dioxygen has been investigated in detail and is indicated by reactions 24-31 (59). For a review discussing the formation 1944
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and fate of peroxyl radicals in aqueous solutions see ref 60. •
CH2C(CH3)2OH + O2 f •OOCH2C(CH3)2OH (24)
2•OOCH2C(CH3)2OH f [HOC(CH3)2CH2OO]2 (25) [HOC(CH3)2CH2OO]2 f O2 + HOC(CH3)2CH2OH + HOC(CH3)2CHO (26) [HOC(CH3)2CH2OO]2 f H2O2 + 2HOC(CH3)2CHO (27) [HOC(CH3)2CH2OO]2 f O2 + 2CH2O + 2•C(CH3)2OH (28) [HOC(CH3)2CH2OO]2 f O2 + HO(CH3)2CCH2OOCH2C(CH3)2OH (29) •
C(CH3)2OH + O2 f •OOC(CH3)2OH
(30)
OOC(CH3)2OH f (CH3)2C)O + HO2•
(31)
•
Reaction 24 is close to diffusion-controlled [k ) 1.8 × 109 M-1 s-1 (40)], and so is the bimolecular decay of the peroxyl radicals which proceeds via a very short-lived (at room temperature) tetroxide [reaction 25, 2k ) (8 ( 2) × 108 M-1 s-1 (59)]. The tetroxide decays by four processes (reactions 26-29), i.e., two molecular routes, the Russell-type reaction 26 and the Bennett-type reaction 27, the formation of the peroxide in reaction 29, and one route that proceeds via free radicals (reactions 28 followed by 30 and 31). The relative importance of the reactions 26-28 has been deduced in a radiolytic study (59) and is reflected in the radiation-chemical yields of certain key products: isobutanediol, cf. reaction 26, relative importance ) unity; hydroxyisobutyraldehyde, cf. reactions 26 and 27, relative importance of reaction 27, ) 1.5; formaldehyde, cf. reaction 28, relative importance of reaction 29, ) 1.3. The relative importance of the formation of the organic peroxide (reaction 29) appeared to be similar to that of reaction 26 (59). Formaldehyde, easy to determine, can be used to monitor •OH formation in such processes where its production is mechanistically possible. In the radiolysis of N2O/O2 (4:1)-saturated aqueous solutions of tertiary butanol the ratio of CH2O produced to •OH radicals expended is about 1/4 (59). An estimate, along these lines, of the amount of •OH radicals produced is possible where there is no ozone present. However, the ozonolytic situation is more complex, and the reactions of the tertiary butanol radical •CH2C(CH3)2OH (reaction 32) and the corresponding peroxyl radical •OOCH C(CH ) OH with ozone (reaction 33) must be taken 2 3 2 into account. •
CH2C(CH3)2OH + O3 f •OCH2C(CH3)2OH + O2 (32)
•
OOCH2C(CH3)2OH + O3 f •OCH2C(CH3)2OH + 2O2 (33)
Under the present conditions, the initial value of the ratio [O3]/[O2] ≈ 0.25 and the rate constant of reaction 32 is expected to be close to diffusion-controlled [benzyl, gas phase: k ) 2.8 × 1010 M-1 s-1 (61); •CH2C(O)O-, aqueous solution: k ) 1.5 × 109 M-1 s-1 (62)]. Peroxyl radicals also react with ozone [reaction 33, k ) 1.8 × 104 M-1 s-1 (63)]. It can be shown that at the above ozone/oxygen ratio, in the order of 10% of the tertiary butanol radicals produced should be transformed into •OCH2C(CH3)2OH. Such oxyl radicals in water undergo rearrangement (reaction 34), a very fast process with rate constants in the order of 107 s-1 (64). This
is in competition with β-fragmentation (reaction 35) (65). The latter process is accelerated in aqueous solution [cf. the tertiary butoxyl radical where the rate constant in the gas phase is in the order of 103 s-1 (65), but it is solvent-dependent (66) and in water is in the order of 106 s-1 (67)]. The presence of the hydroxyl group in the position β to the oxyl function has an accelerating effect on this decay (68). The radical • OCH2C(CH3)2OH decays particularly fast (reaction 35), k ) 1.4 × 107 s-1 in the gas phase (65); this value may be taken as a lower limit to its aqueous-phase value. For the same reason, reaction 34 may also be accelerated compared with the decomposition of tertiary butoxyl. It is thus seen that the quantum yield of formaldehyde depends, among other things, also on the value of the branching ratio k34/k35 whose upper limit is not expected to exceed unity; it may be smaller if an accelerating effect of the aqueous medium operates on k35, as it does in the case of tertiary butoxyl but from a much lower base. •
OCH2C(CH3)2OH f •CH(OH)C(CH3)2OH
(34)
OCH2C(CH3)2OH f CH2O + •C(CH3)2OH
(35)
•
The radical •CH(OH)C(CH3)2OH, having added dioxygen to give rise to a second-generation peroxyl radical, can undergo O2•- elimination similar to reaction 31, either as such (this reaction may however be relatively slow if the behavior of the corresponding radical •OOCH(OH)CH3 derived from ethanol whose rate constant is 52 s-1 (69), vs > 106 s-1 in dihydroxyperoxyl radicals such as •OOCH(OH)2 (70), is any guide), or after a further cycle of oxidation, peroxylation, rearrangement, and HO2• formation (reactions 36-43). •
CH(OH)C(CH3)2OH + O2 f •OOCH(OH)C(CH3)2OH (36)
•
CH(OH)C(CH3)2OH + O3 f
I + II f O2 + HOC(CH3) 2CH2OH + HOC(CH3)2COOH (44) f O2 + 2HOC(CH3)2CHO + H2O
(45)
f H2O2 + HOC(CH3)2CHO + HOC(CH3)2COOH (46) f O2 + •OCH2C(CH3)2OH + •
OCH(OH)C(CH3)2OH (47)
f O2 + HO(CH3)2CCH2OOCH(OH)C(CH3)2OH (48) II + II f O2 + HOC(CH3)2CHO + HOC(CH3)2COOH + H2O (49) f O2 + 2•OCH(OH)C(CH3)2OH
(50)
The extent to which the second- and third-generation peroxyl radical reactions contribute in the tertiary butanol/ ozone system depends, among other things, primarily on the branching ratio k34/k35. Acids are expected to be produced, e.g. in reactions 39, 43, 44, 49, and 52. The experiments done near neutral pH immediately show that this is indeed the case. In particular, the discovery of hydroxyisobutyric and formic acids (Table 1) implies that reaction 34 which provides a channel to the more highly oxidized products is significant. The presence of a small amount of acetic acid derives from the reaction of the 2-hydroxyprop-2-yl radical with ozone (reactions 51-54). •
C(CH3)2OH + O3 f •OC(CH3)2OH + O2
(51)
•
•
OCH(OH)C(CH3)2OH + O2 (37)
•
owing to the smaller concentrations of II and especially III compared to I.
•
OCH(OH)C(CH3)2OH f C(OH)2C(CH3)2OH (38)
•
OCH(OH)C(CH3)2OH f HCO2H + •C(CH3)2OH
•
OOCH(OH)C(CH3)2OH f HC(O)C(CH3)2OH + HO2• (40)
(39)
•
OOCH(OH)C(CH3)2OH + O3 f •
OCH(OH)C(CH3)2OH + 2O2 (41)
•
C(OH)2C(CH3)2OH + O2 f •OOC(OH)2C(CH3)2OH (42) OOC(OH)2C(CH3)2OH f HOC(CH3)2CO2H + HO2• (43)
•
The relatively long-lived (with respect to HO2•/O2•elimination) second-generation peroxyl radical •OOCH(OH)C(CH3)2OH (II, the first-generation peroxyl radical obviously being •OOCH2C(CH3)2OH I) is expected also to take part in various termination reactions, analogous to reactions 25 followed by 26-29, in contrast to the third-generation peroxyl radical •OOC(OH)2C(CH3)2OH (III) which should eliminate HO2• very fast, cf. ref 70. These reactions whose branching ratios are unknown are expected also to give rise to aldehydes and acids. Further peroxyl-radical termination reactions exist, among them reactions 44-48, which are probably of minor importance compared to the self-termination of peroxyl radical I,
OC(CH3)2OH f CH3CO2H + •CH3
(52)
•
CH3 + O2 f CH3OO•
(53)
CH3OO• + RO2• f f CH2O, other products
(54)
A series of experiments has been carried out at neutral pH, where the HO2• radical is replaced by the O2•- radical. It has been observed that in the presence of millimolar and higher concentrations of tertiary butanol this leaves the rate of photolytic ozone decomposition practically unaffected compared with the situation at pH 2 (modeling indicates the same). In the absence of an •OH-scavenging cosolute, however, the ozone decomposition rate is much faster than at pH 2 since O2•- is the more effective chain carrier compared with HO2•. Also under these pH conditions, the OH--induced spontaneous ozone decomposition does not yet interfere since the half-time of this process, in the order of 10 h in the presence of 0.01 M tertiary butanol, see above [in the absence of an organic cosolute it is much shorter, cf. ref 49], is long compared with the duration of the experiment. It is important here that in contrast to the acidified solutions, neutral solutions may be analyzed by gas chromatography and so the consumption of tertiary butanol be determined as well as (by ion chromatography) the production of acids. Thus at an ozone concentration of 3 × 10-4 M and tertiary butanol of 1 × 10-4 M, about half of the tertiary butanol has reacted when the ozone has disappeared, i.e., the efficacy ∆[t-BuOH]/ ∆[O3] ≈ 0.17 (Table 1) under these conditions. At high conversion, protection of the substrate by the products may partly explain low η values at low initial tertiary butanol concentrations. If the initial concentration of tertiary butanol VOL. 37, NO. 9, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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is higher, the efficacy is better (≈0.26 at 1 mM) because a lesser part of •OH is lost by reactions with O3 and with the tertiary-butanol-derived products which are largely more reactive toward •OH than tertiary butanol itself (16). The pH drop corresponds to the production of acid in the order of 10-5 M. Ion chromatography shows the formation of formic, 2-hydroxyisobutyric, and acetic acids (Table 1, average of three or four values that display considerable scatter in the order of 50%). The fact that the former two are produced even at relatively high tertiary butanol concentrations rules out their formation as secondary products via the corresponding aldehydes. An aldehydic product of uncertain identity was observed whose chromatographic behavior matched that of acetaldehyde, in amounts of less than a tenth of those of hydroxyisobutyraldehyde. The value of Φ(-O3) ≈ 0.64 in the presence of millimolar or higher concentrations of tertiary butanol indicates that the quantum yield of ozone destruction due to direct photolysis (reactions 3-6) must be lower than this value. This means that in the condensed phase, these reactions are to a greater or lesser extent reversed in the solvent cage (reactions 55 and 10), in competition with the reaction of O(1D) with water (reaction 7).
O(1D) + O2(1∆g) f O3
(55)
The attempt has been made by kinetic modeling on the basis of the reaction mechanism above, to estimate the yield of the •OH photolytically formed from the aqueous ozone (reactions 7 followed by 8). To image the experimental value η(-BuOH) ) 0.26 obtained at a tertiary butanol concentration of 10-3 M (Table 1), a quantum yield of about 0.10 must be assumed for •OH formed in reaction 7 followed by reaction 8. For a tertiary butanol concentration set at 10-4 M, this code gives η(-BuOH) ) 0.20, vs an experimental value of 0.17 (Table 1). For 10-3 M tertiary butanol, one obtains η(CH2O) ) 0.08 vs an experimental value of 0.07 (Table 1). For a hypothetical situation where the tertiary-butanol radical is assumed to react neither with O2 nor with O3, the quantum yield of ozone disappearance is calculated at 0.52; this is through direct photolysis only, in the hypothetical absence of organic free-radical participation. This result suggests that the free radicals generated from tertiary butanol and their descendants promote the rate of ozone decomposition by about 20%, to the observed value of 0.64. Further, for another hypothetical situation with only the reactions of ozone with free radicals, such as reactions 32, 33, 37, 41, and 51 set at zero, the quantum yield of ozone disappearance would be about 0.55, and η(CH2O) ) 0.06, i.e., 25% lower than it is in the presence of these reactions, while η(-BuOH) comes out only slightly lower, by about 4%. This points to the mechanistic significance in this system of the reactions of ozone with various organic free radicals. It is felt that the use of this modeling procedure where the direct photolytic ozone decomposition step, whose actual kinetics varies from zeroth to first order, was approximated by a first-order reaction is adequate (even though somewhat rough) in view of the limited accuracy of the experimental data; also, the concentration gradient in the rate of generation of the free radicals across the sample was neglected. The reactions and rate constants used are referred to in Table 2, the rate of initiation being determined by the light intensity and a quantum yield of ∼0.6 of reactions 3 plus 4 (reactions 5 plus 6 are offset by ozone restitution via reaction 10). A value of about 0.52 for the quantum yield of direct ozone photocleavage and of about 0.10 for direct OH radical generation [i.e. Φ(O + H2O f 2•OH) ∼ 0.05] leaves a value of nearly 0.5 for the quantum yield of H2O2 formation via reaction 7. Peroxyl radical-termination followed by reactions such as reaction 27 give rise to further H2O2. Our experimental 1946
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TABLE 2. Rate Constant Values Used in the Modeling Procedure reaction 3+4 5+6 7 8, 9 10 11 12 16 17 23 24, 25, 30, 36, 42, 53, 54 26 27 28 29 31 33, 41 34 35 38 39 40 43 44, 45, 46, 47, 48 49, 50 52 •OH + H O 2 2 O2•-+ HO2• • HO2 + HO2• O(3P) + (CH3)3COH a
rate constant (M-1 s-1) 1.8 × 10-3 (s-1), k5+6 ) 0.1 k3+4 1.8 × 10-4 (s-1), k5+6 ) 0.1 k3+4 1.8 × 1010 (26) k8/k9 ) 0.11a 4 × 109 (31) 5 × 106 (32) 1.1 × 108 (16) 1.6 × 109 (1) 1.2 × 106 (32) 6 × 108 (16) 2 × 109 assumed; cf. ref 60
k26/(k26 + k27 + k28 + k29) ) 0.2; cf. ref 59 k27/(k26 + k27 + k28 + k29) ) 0.3; cf. ref 59 k28/(k26 + k27 + k28 + k29) ) 0.25; cf. ref 59 k29/(k26 + k27 + k28 + k29) ) 0.25; cf. ref 59 6 × 102 (s-1) (69) 1 × 104 (63) 1 × 107 (s-1) assumed, see text 1.4 × 107 (s-1) (65), see text, cf. ref 64 1 × 107 (s-1) cf. ref 64 1 × 107 (s-1) assumed, cf. k38 5 × 101 (s-1) assumed, cf. ref 69 1 × 106 (s-1) assumed, cf. ref 70 2 × 109 assumed, cf. ref 60; k44 ) k45 ) k46 ) k47 ) k48 ) 1/5 ktotal 2 × 109 assumed, cf. ref 60; k49 ) k50 ) 1/2 ktotal 1 × 106 (s-1) assumed, cf. ref 67 2.7 × 107 (16) 9.7 × 107 (52) 8.3 × 105 (52) 2.5 × 106 assumed, cf. ref 36b
This ratio leads to η(- t-BuOH) ) 0.26 at [t-BuOH] ) 10-3 M, see text.
finding is η(H2O2) at ∼0.9 in the presence of 0.01 M tertiary butanol. At the H2O2 concentrations that are built up under these conditions, its destruction by reaction with ozone is insignificant even when the solutions are neutral since the pKa of H2O2 is 11.6 and the rate constant of ozone with HO2at 5.5 × 106 M-1 s-1 (1) is not very large. Also at millimolar tertiary butanol concentrations and above, the formation of HO2•/O2•- via the reaction of •OH with the H2O2 being built up during the irradiation is negligible. The Efficiency of Ozone/UV as an AOP. It has been shown above that the quantum yield of direct ozone photolysis is Φ ≈ 0.5. Only about 10% of the photolyzed ozone furnish •OH (O + hν + H O f 2•OH + O ), i.e., the quantum yield 3 2 2 of •OH production is relatively low, Φ ≈ 0.1, when compared with the photolysis of H2O2 [Φ(•OH) ) 1.0 (71)]. Thus, neither UV quanta nor ozone are efficiently converted into •OH. The pollutant to be degraded must compete for •OH with water constituents that are omnipresent such as bicarbonate [k ) 8.5 × 106 M-1 s-1 (16)] and carbonate [k ) 3.9 × 108 M-1 s-1 (16)] as well as organic matter (NOM). Rate constants (in units of mgC L-1 s-1) center around 2 × 104 [(1.3-2.5) × 104 (72), data evaluated by G. R. Peyton; (2.3 ( 0.7) × 104 (73); 1.7 × 104 (74), 1.6 × 104 (75); 2.3 × 104 Nowell and Hoigne´ cited in ref 75; 1.9 × 104 (76); (1.9-2.7) × 104 (77); ∼1 × 104 (pulse radiolysis of a brown-water from a mountain lake in the Black Forest, M. N. Schuchmann and C. von Sonntag, unpublished results). The reaction of •OH with the ozone itself [k ) 1 × 108 M-1 s-1 (16)] must also be taken into account and gains in importance as the ozone concentration is raised. As the reaction proceeds, H2O2 builds up due to the major pathway of ozone photolysis (Φ(H2O2) ≈ 0.5) as well as in the course of the peroxyl radical reactions (an example is given
above). However, H2O2 has only a relatively low reactivity toward •OH [k ) 2.7 × 107 M-1 s-1 (16)]. In any case in this reaction, an HO2•/O2•- radical is formed which with ozone regenerates •OH. Thus in the UV/O3 system this reaction does not have as negative an influence as it has in another AOP, i.e., the photolysis of H2O2, where higher H2O2 concentrations are typically used than those that build up in the UV/O3 process, and where a regeneration of •OH is absent. Despite the fact that the UV/O3 process has such a low efficiency for •OH formation, it may under certain conditions be favorable compared with the corresponding UV/H2O2 process because of the much higher (165-fold at 254 nm) optical absorption coefficient of ozone as compared with H2O2. Moreover, and in contrast to H2O2, any residual ozone will decay in the water pipeline and not reach the consumer. The UV/O3 process has been shown to be effective in the elimination of chlorinated hydrocarbons (7-9), i.e., under conditions where the pollutant neither reacts with ozone nor absorbs much of the UV light [cf. ref 78]. These are the adequate conditions for its application. Nevertheless, there is still a steady output of studies that investigate the potential usefulness of the UV/O3 process with systems that react already by themselves readily with ozone. A case in point is phenol (79). Phenol reacts with ozone with a rate constant of 1.3 × 103 M-1 s-1 and with phenolate with 1.4 × 109 M-1 s-1 (1). The pKa value of phenol is 10.0, and it is easy to show that hence the rate of reaction of phenol with ozone at pH 7 is k ≈ 106 M-1 s-1. At a phenol concentration of 1 × 10-3 M (79) the rate of ozone degradation is thus near 103 s-1. For the UV/O3 process to become effective, the photolysis of ozone must compete favorably with the thermal ozone reaction. A typical example of the rate of ozone photolysis using the low-pressure mercury lamp as the UV-light source is shown in Figure 1 where the observed first-order rate constant is only on the order of 10-2 s-1. Using stronger light sources, the medium-pressure Hg arc instead of the lowpressure Hg arc on which these calculations are based would not be able to compensate for this order-of-magnitude difference in rate between the dark reaction and the photolysis. Thus, any procedures that use the UV/O3 system with ozone-reactive pollutants are unlikely to yield better results than ozonation alone. It therefore appears to be more cost-effective in such cases first to carry out the ozone treatment without irradiation, followed by the O3/UV process to deal with any ozone-refractory products of the ozonolysis (11).
Acknowledgments We would like to thank a referee for several most helpful suggestions and Dr. G. R. Peyton for valuable discussions.
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Received for review September 25, 2001. Accepted February 5, 2003. ES0113100
