Photocatalytic transformations of CCl3Br, CBr3F, CHCl2Br and CH2BrCl in aerobic and anaerobic conditions

Applied Catalysis B: Environmental 29 (2001) 23–34

Photocatalytic transformations of CCl3Br, CBr3F, CHCl2 Br and CH2 BrCl in aerobic and anaerobic conditions Paola Calza a , Claudio Minero a , Anastasia Hiskia b , Elias Papaconstantinou b , Ezio Pelizzetti a,∗ a

Dipartimento di Chimica Analitica, Università di Torino, 10125 Torino, Italy Institute of Physical Chemistry, NCSR Demokritos, 15310 Athens, Greece

b

Received 30 January 2000; received in revised form 2 June 2000; accepted 9 June 2000

Abstract Phototransformations of halomethanes containing chlorine and bromine (CCl3 Br, CHCl2 Br, CH2 ClBr) or bromine and fluorine (CBr3 F) have been investigated under aerobic and anaerobic conditions both in homogeneous system and heterogeneous photocatalysis. For all of those compounds, the complete disappearance of the primary compound and the stoichiometric concentration of halides was achieved. Several halogenated intermediates and oxygenated compounds were identified, so that it was possible to predict the degradation pathways followed by such halomethanes. Whereas the reductive steps are predominant in the initial degradation of CCl3 Br and CBr3 F, the oxidative steps are predominant in the initial CH2 ClBr steps. The two pathways have comparable importance for CHCl2 Br degradation. Methanol, acting as a hole scavenger, strongly increases the rate of disappearance for CCl3 Br and CBr3 F. © 2001 Elsevier Science B.V. All rights reserved. Keywords: Photocatalysis; Titanium dioxide; Halomethanes

1. Introduction The release of bromine containing organic compounds into the atmosphere is of interest because of their potential adverse impact on stratospheric ozone [1]. Species such as CH2 ClBr, CHBrCl2 , CHBr2 Cl and CHBr3 are observed, all of which appear to have natural oceanic sources [2–5]. CH2 ClBr is a natural gas, formed by algal biological processes, and hence delivers bromine into the earth’s atmosphere. Although bromine compounds are present in the atmosphere in much smaller quantity than their ∗ Corresponding author. Tel.: +39-011-6707630; fax: +39-011-6707615. E-mail address: [email protected] (E. Pelizzetti).

CFC equivalents, bromine leads to a much higher destruction of ozone than chlorine [6,7]. In addition, such compounds are found to be not only ubiquitous contaminants in the atmosphere, but also in the subsurface aquifers. Their degradation mechanism can give information about their environmental fate, and lead to the improvement of methods of remediation. Photocatalysis over irradiated TiO2 has been shown to dechlorinate the more heavily substituted homologues of chloromethane derivatives [8–10] and bromomethane derivatives as well as other chlorinated ethanes and ethenes [11,12]. The photocatalytic transformation of bromochloromethanes and tribromofluoromethane is presently being investigated in aerobic and anaerobic conditions.

0926-3373/01/$ – see front matter © 2001 Elsevier Science B.V. All rights reserved. PII: S 0 9 2 6 - 3 3 7 3 ( 0 0 ) 0 0 1 8 8 - 0

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2. Experimental 2.1. Material and reagents All degradation experiments were carried out using TiO2 Degussa P25 as photocatalyst. The TiO2 powder was cleaned as previously reported [13]. Trichlorobromomethane (Aldrich), dichlorobromomethane (Aldrich), bromochloromethane (Aldrich), tribromofluoromethane (Aldrich), formic acid (Carlo Erba), formaldehyde (Merck), methanol (Merck) and tert-butanol (Aldrich) were used as received. Sodium chloride (Merck), potassium bromide (Carlo Erba) and sodium fluoride (Merck) were used after drying. 2.2. Irradiation procedures Irradiations were carried out on 5 ml of suspension containing 10 mg l−1 halomethane and 200 mg l−1 of TiO2 , using a 1500 Watt Xenon lamp (Solarbox, CO. FO. MEGRA, Milan, Italy) simulating AM1 solar light and equipped with a 340 nm cut-off filter. The irradiation spectrum and the cells have been described elsewhere [14,15]. The total photonic flux (340–400 nm) in the cell and the temperature have been kept constant for all the experiments. They were 1.35×10−5 einstein min−1 and 50◦ C, respectively. Experiments were run at the natural pH of the aqueous suspension, at pH 4 after adjustment with HNO3 and at pH 10 and 12 after addition of NaOH. The experiments in absence of oxygen were prepared by purging the irradiation cells filled with the TiO2 suspension with helium for 30 min. Following this procedure, the required volume of substrate stock solution was injected into the cell.

have already been reported [13]. Either GC Varian STAR 3400 with a FID detector or a GC (HP 6890) equipped with a mass detector (HP 5973, operating in EI ionization mode) was utilized. Primary compound and the halogenated intermediates formed were identified by their EI mass spectra and on the retention time equivalence with standards. Under the reported conditions [13], the retention times were 4.45, 5.76, 5.26, 7.03, 8.28, 8.47 min for CH2 Cl2 , CHCl3 , CH2 ClBr, CHCl2 Br, CCl3 Br and CBr3 F, respectively. Formaldehyde was derivatized as previously described [16]. The chromatographic analysis was accomplished by HPLC using a Rheodyne injector, a RP C18 column (Lichrochart, Merck, 12.5 cm×0.4 cm, 5 ␮m packing), a high pressure two-pumps gradient (Merck Hitachi L-6200 and L-6000 pumps) and UV–VIS detection (Merck Hitachi L-4200). Formic acid was detected using ion chromatography. A Biotronik IC 5000 apparatus equipped with a Biotronik BT III OS type anion exclusion column was used. The elution was performed using H2 SO4 5×10−4 M at 1 ml min−1 flow rate. The same apparatus was used for bromide and chloride analysis. Suppressed ion chromatography was used with conductivity detection. The separation was carried out on a BT1AN column (20 cm length, 4 mm i.d., Biotronik) using an alkaline buffer eluent containing NaHCO3 (3 mM) at the flow rate of 1.5 ml min−1 ; under these conditions, the retention time of bromide and chloride were 5.5 and 4.3 min, respectively.

3. Results and discussion 3.1. Phototransformation of trichlorobromomethane and tribromofluoromethane

2.3. Analytical procedures The contents of the cell were filtered through a 0.45 ␮m cellulose acetate filter (Millipore HA) and analyzed by the appropriate analytical technique. The disappearance of the primary compound was followed using a purge and trap system (Tekmar LSC 2000) equipped with a Vocarb 3000 trap and a cryofocusing module connected with a 60 m DB5 column (Supelco, 0.32 mm i.d., 0.25 ␮m coating). The purge and trap’s and GC operative parameters

3.1.1. Homogeneous photolysis While CBr4 is photochemically unstable and can be easily photocatalytically degraded [17], CCl4 is photochemically very stable [13]. This suggests that the presence of different halogens in the same molecule will result in different photolytic stability. In this study, experiments were carried out on molecules containing chlorine and bromine (CCl3 Br) and bromine and fluorine (CBr3 F) to evaluate the stability under thermal and photolytic conditions.

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Fig. 1. (A) Disappearance of trichlorobromomethane (CO =5×10−5 M) in the dark and in the light (photolysis) under aerobic and anaerobic condition; (B) disappearance of trichlorobromomethane (CO =5×10−5 M) for direct photolysis at different pH.

The disappearance of CCl3 Br in the dark and for direct photolysis, under aerobic and anaerobic conditions, is reported in Fig. 1. In the dark, air has no significant effect and, after 4 h, a loss of 15–20% of CCl3 Br was observed. Under illumination after 4 h 80% of primary compound disappears. The homolysis of the C–X bond to halogen atoms and haloalkyl radicals [18–20] occurs according to Eq. (1)

The decomposition of halogenated hydrocarbons is carried out either by oxidative species or conduction band electrons. CCl3 Br and CBr3 F contain a carbon atom in its higher oxidation state. Therefore, their photocatalytic transformations in aqueous solution are unlikely to be initiated by an oxidative attack. They can react with the electrons of the conduction band according to

CCl3 Br + hν → • CCl3 + • Br

− → • CXa Yb−1 + Y− CXa Yb + eCB

(1)

while the presence of oxygen does not significantly influence the rate of CCl3 Br degradation, different pH conditions bring different rates of disappearance. The rate is highest around pH 5.5, and decreases while raising or lowering the pH: 2.0×10−5 M h−1 at pH 4, 3.3×10−5 M h−1 at pH 5.5, 1.8×10−5 M h−1 at pH 10 and 1.2×10−5 M h−1 at pH 12. Bromide evolution rate is three times faster than [Cl− ]/3, as shown in Fig. 1. This is expected on the basis of C–Br and C–Cl bond homolysis [18–20]. In the case of CBr3 F, losses were found neither in the dark nor under illumination under the conditions and observation times reported above. The presence of fluorine gives photolytic stability to the molecule. 3.1.2. Photocatalytic transformations under helium or air Photoexcitation of titanium dioxide with light below 380 nm generates oxidative valence band holes (E0 =+3.1 V versus NHE at pH=0) and reductive conduction band electrons (E0 =−0.1 V): − + h+ TiO2 + hν → eCB VB

(2)

(3)

The photocatalytic degradations in the presence of TiO2 were followed by the disappearance of the primary compound, the production of halide and the formation of the intermediate as a function of time. The disappearance of the primary compound, under aerobic and anaerobic conditions, together with the evolution of halides are shown in Fig. 2 for CCl3 Br

Fig. 2. Photocatalytic degradation of 5×10−5 M CCl3 Br on TiO2 200 mg l−1 at pH 5.5 and the production of bromide and chloride ions under aerobic and anaerobic conditions.

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≡Ti3+ –O2− –H + O2 → ≡Ti4+ –O2− –H + O2 •− ↔ HO2 •

Fig. 3. Photocatalytic degradation of 5×10−5 M CBr3 F on TiO2 200 mg l−1 at pH 5.5 and the production of bromide and fluoride ions under aerobic and anaerobic conditions.

and in Fig. 3 for CBr3 F, while the initial degradation rates of CCl3 Br, CBr3 F and, for comparison, CCl4 and CBr4 , are reported in Table 1. Similar steps of degradation occur for all compounds and the reductive pathways are initially predominant; particularly, in the oxygen-free system. Nevertheless, the initial rates for the halogenated compounds are different. Under helium, CBr4 is three times faster than CCl4 , while CBr3 F is 10 times slower than CBr4 , the presence of the C–F bond strongly stabilizing the halomethane.The presence of oxygen competes with − generating the less reactive O2 −• /HO2 • [21] eCB − ≡Ti4+ –O2− –H + eCB → ≡Ti3+ –O2− –H

(4)

Table 1 Initial degradation rate of halomethanes (initial concentration CBr4 1.5×10−5 M, CCl4 1.5×10−5 M, CCl3 Br 5×10−5 M, CBr3 F 3.67×10−5 M) on TiO2 200 mg l−1 under air or helium, or helium and methanol (0.01 M), and the corresponding initial rate of halide production Rate (M h−1 )×105

Air

Helium

He+CH3 OH

CCl4 (19) Cl− CBr4 (15) Br− CCl3 Br Cl− Br− CBr3 F Br− F−

5.6 21.0 5.4 10.2 7.1 4.9 2.6 1.0 2.8 1.0

24.0 20.0 78.0 72.0 11.7 7.1 4.3 7.3 6.6 2.2

42.0 – – – 26.7 18.6 6.9 25.0 26.4 3.7

(pKa = 4.8)

(5)

thus decreasing the weight of the reductive pathways. In the absence of oxygen, all of the degradation rates are increased. Nevertheless, oxygen has a small effect on CCl3 Br, while the effect is remarkable on CCl4 , CBr3 F and CBr4 . In particular, under aerobic conditions, the rate of CCl4 is close to that of CBr4 ; while the rate of CCl3 Br is almost twice that of CCl4 , and the rate of CBr3 F is 1/5 of CBr4 . The formation of chloride and bromide ions is also reported in Table 1. Aliphatic halogen compounds are − quantitatively dehalogenated by reaction with eaq − eaq + RX → RX− → R• + X−

(6)

On the basis of data reported in the literature [22,23], the rate of halide release is expected to be in the order: Br>Cl>F. In fact, for CCl3 Br the initial rate of formation of Br and Cl/3 is 2:1, while for CBr3 F is Br/3: F∼1:1. Halide evolution depends on initial reductive attack and on the hydrolysis of the radical intermediates. In the oxygen-free system, the rate d[X− ]/dt is similar to –d[CX]/dt for the four compounds, while in air the rate d[X− ]/dt is 2–4 times that of −d[CX]/dt for CCl4 , CBr3 F and CBr4 , and some only for CCl3 Br. With helium and methanol, halide evolution has the same behaviour as above, the only difference being an increase in the rate of evolution. The halide evolution from CCl3 Br under helium shows Br− and Cl− /3 initially close to each other, while at longer times Br− evolution becomes slightly faster than Cl− /3. In air, Br− evolution is slightly faster than Cl− /3 (e.g. Cl− 3.6×10−5 M h−1 and Br− 1.6×10−5 M h−1 ). Under helium, Br− and Cl− faster approach the stoichiometric value; after 4 h of irradiation, 85 and 66% of the expected bromide and chloride concentrations are reached, respectively, while in air 80 and 50% of expected Br− and Cl− , respectively, are formed after the same irradiation time. The kinetic of the process is complex; while initially a first-order kinetics is followed, and Cl− and Br− are released with the same ratio both in air and He, after 1 h the kinetics changes and the bromide is released more faster. Also after 1 h, CCl3 Br almost completely disappears and the change in kinetics is believed to be probably due

P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34

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Fig. 4. Influence of methanol (0.01 M) on the photocatalytic degradation of CCl3 Br (A) and CBr3 F (B) on TiO2 200 mg l−1 at pH 5.5; disappearance of the primary compound and halide production under anaerobic conditions.

to the halogenated intermediates, especially the chlorinated compounds which probably release halide at different rates. Similarly, with CBr3 F, a first-order kinetics followed during the first 30 min, which changes if the irradiation time is prolonged; bromine and fluorine are released at comparable rates both in aerobic and anaerobic conditions, although in the early part of the reaction, fluoride shows an induction period. 3.1.3. Effect of methanol on CCl3 Br and CBr3 F degradation rates The influence of an electron donor on CCl3 Br and CBr3 F was studied under anaerobic conditions and the obtained results are shown in Fig. 4. Similar to that observed for CCl4 [21] and CBr4 [17], methanol acts as a hole scavenger and results in an increase in degradation rates (see Table 1). Methanol reacts with the photoproduced holes (or • OHads ) [24] CH3 OH + • OHads → • CH2 OH + H2 O

(7)

enhancing the available electrons. The reduction potential of • CH2 OH (−0.74 V) [25] is negative enough to inject electrons to the conduction band of TiO2 (doubling effect) [26]: • CH

2 OH

+ TiO2 → CH2 O + H+ + TiO2 (eCB )

(8)

or directly reduce CX3 Y as follows CX3 Y + • CH2 OH → HCHO + • CX3 (or • CX2 Y) + H+ + Y− (or X− )

(9)

Under helium and methanol, for CCl4 the rate of disappearance is almost twice. For CBr3 F and CCl3 Br

the rate is almost three times; this could support the hypothesis that • CH2 OH is contributing to the reaction progress. As expected, the presence of methanol causes an increase in the rate of disappearance of CCl3 Br (t1/2 =5 min); Cl− and Br− ratio is close to 1:1 and after 1 h of irradiation, 85 and 66% of stoichiometric bromine and chlorine appear, respectively. In CCl3 Br degradation some halogenated intermediates (CH2 ClBr, CHCl3 and CH2 Cl2 ) have been identified. Dichloromethane is the main intermediate, while bromochloromethane forms rapidly and disappears quickly. The presence of methanol also has a positive effect on the rate of degradation of CBr3 F. In fact, under these experimental conditions CBr3 F is the compound with the highest degradation rate, whereas it is the one with the lowest degradation rate in the presence of air. With He and methanol, fluoride shows a more pronounced induction period with Br− /3 appearing much faster than F− ; when CBr3 F has virtually disappeared, only 5% of fluoride ions are formed. 3.1.4. Phocatalytic degradation mechanism of trichlorobromomethane The intermediates observed in CCl3 Br degradation under aerobic and anaerobic conditions are given in Fig. 5. As previously shown, bromide ion is released according to the following reaction − → • CCl3 + Br − CCl3 Br + eCB • CCl , 3

(10)

reacting with • OH radicals, can produce HOCCl3 , a short-lived intermediate. The sequence

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Fig. 5. Photocatalytic degradation of 5×10−5 M CCl3 Br on TiO2 200 mg l−1 at pH 5.5; intermediate formation under aerobic (A) and anaerobic (B) conditions.

from HOCCl3 to CO2 and chloride is fast [27] HOCCl3 → COCl2 + H+ + Cl−

(11)

COCl2 + H2 O → CO2 + 2Cl− + 2H+

(12)

• CCl 3

can also be reduced to give CHCl3

• CCl

+ e− + H+ → CHCl3

3

(13)

Further reduction of chloroform (observed as an intermediate) leads to • CHCl2 radical and then to dichloromethane. Even if the formation of bromide ions is favoured [22,23], chloride ions are also formed. In the initial phase this may imply that, besides the formation through Eqs. (10) and (11), the chloride can also be produced in the formation of • CCl2 Br radical − → • CCl2 Br + Cl− CCl3 Br + eCB

(14)

This is particularly evident in the absence of oxygen. In fact, the analysis of the intermediates obtained within half an hour, both halogens are released with comparable rates. This is confirmed by the intermediates which have been identified. Among the species initially formed is CH2 ClBr, which must originate from the • CCl2 Br radical via the sequence • CCl

2 Br

+ e− + H+ → CHCl2 Br −

(16)

+ e + H → CH2 ClBr

(17)

CHCl2 Br + e → • CHClBr + Cl • CHClBr

(15)

−

−

+

The intermediates detected during CCl3 Br degradation were CHCl3 , CH2 Cl2 , CH2 ClBr, HCOOH and HCHO. Their profiles are shown in Fig. 5a. CCl3 Br

generates more CH2 Cl2 than CHCl3 , while among the oxygenated intermediates, formaldehyde is formed in a higher concentration than HCOOH. The presence of CH2 Cl2 in high concentration can be justified because it may be formed via two routes: from CHCl3 reduction and from CHCl2 Br loss of bromide. After 30 min, the main intermediates are CH2 Cl2 and HCHO, that together account for 15% of the CCl3 Br disappeared. CH2 ClBr accounts for 7%, while CHCl3 and HCOOH account for 7 and 5%, respectively. At longer irradiation times, the halogenated compounds disappear and oxygenated species, such as formic acid and formaldehyde, are formed. In particular, formaldehyde becomes the main intermediate accounting for 35%. Formaldehyde can be produced during short irradiation times, by dechlorination and debromination of CH2 ClBr, through the overall stoichiometric reaction: CH2 ClBr + H2 O → HCHO + HCl + HBr

(18)

while for longer irradiation time, HCHO can be generated through oxido-reductive sequence from the degradation of CH2 Cl2 . In fact, it is well known that under anaerobic conditions, HCHO is the main product obtained from CH2 Cl2 [13] degradation. On the basis of the intermediates identified in the different experimental conditions, the mechanism shown in Scheme 1 is suggested, under which account for all species observed and for possible transient radicals which have been reported. In the photocatalytic degradation of CCl4 , the two-electron pathway via dichlorocarbene formation was shown to play an important role [28]. Under the present conditions, no dimers were detected.

P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34

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Scheme 1. Pathways for the transformation of bromochloromethanes under oxygen-free photocatalytic conditions.

The presence of oxygen modifies the degradation pathways and kinetics. As intermediates, CHCl3 and CH2 Cl2 , resulting from • CCl3 radical although in lower concentration and some oxygenated species are identified, while no species containing chlorine and bromine have been detected, suggesting that under such experimental conditions the Br elimination route is the favoured one. • CCl3 radical can react very fast with oxygen, according to the sequence [29–32] CCl3 + O2 → CCl3 O2 •

(19)

CCl3 O2 • → COCl2 + ClO•

(20)

COCl2 is easily hydrolyzed to CO2 and HCl. The profiles of the production of intermediates are depicted in Fig. 5b. After 1 h, CHCl3 and CH2 Cl2 account for 5 and 7%, respectively. Oxygenated compounds are also formed and HCHO becomes the main intermediate after 30 min. The average oxidation number of the carbon atom nC of the different quantified compounds was calcu-

lated as a function of time. The oxidation number of H, Cl, Br and O were taken as +1, −1, −1 and −2, respectively. nC is calculated through the formula P ci (nC )i nC = P ( ci )i where ci is the concentration of the species in which carbon has the oxidation state (nC )i . Since complete mineralization and a stoichiometric formation of bromide and chloride are achieved, the difference between the values calculated from all the quantified species and the stoichiometric value was attributed to CO2 evolution. This may cause an uncertainty in the nC evaluation. Considering the profiles shown in Fig. 6, it is evident that the reductive pathways are predominant both under aerobic and anaerobic conditions. In absence of oxygen, nC decreases as a function of time and after 4 h reaches minimum at nC ∼ =2.2. In the presence of oxygen after 1 h, nC ∼ =3.2, and then slowly increases

30

P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34 Table 2 Initial degradation rate of halomethanes (initial concentration CHBr3 4.5×10−5 M, CHCl3 8.5×10−5 M, CHCl2 Br 6.1×10−5 M) on TiO2 200 mg l−1 under air or helium, and the corresponding initial rate of halide production Rate (M h−1 )×105 (19)

Fig. 6. The average oxidation number nC calculated for CCl3 Br, CHCl2 Br and CH2 ClBr system as a function of time in the presence and absence of air.

toward +4, corresponding to the total evolution of CO2 . The variation of nC is similar to the one shown by CCl4 and CBr4 , except for the continuous decrease in nC under anaerobic conditions. This may be linked to the slow rate of formation of the halides. An increase in nC may be realized only for longer irradiation times. In order to get further information, the degradation of some bromochloromethanes has been carried out. 3.2. Photocatalytic degradation of dichlorobromomethane Bromodichloromethane was irradiated in the presence of TiO2 under aerobic and anaerobic conditions. The disappearance of the primary compound, together with the production of halide evolution and the intermediates produced in presence of oxygen are shown

CHCl3 Cl− CHBr3 (15) Br− CHCl2 Br Cl− Br−

Air

Helium

6.5 13.0 7.2 13.2 20.6 12.2 10.1

7.0 12.0 6.0 13.2 22.2 10.8 9.2

in Fig. 7a, while the profiles observed in absence of oxygen are reported in Fig. 7b. A comparison between the rates of disappearance of CHCl2 Br, CHCl3 and CHBr3 is summarized in Table 2. Oxygen does not seem to influence the rate of degradation of the parent molecules since their rates are virtually not modified by the presence or absence of oxygen. This suggests that for these halogenated compounds the oxidative and the reductive pathways are kinetically balanced under air-equilibrated conditions. While CHCl3 and CHBr3 degrade with similar initial rates, the photocatalytic degradation of CHCl2 Br occurs at a higher rate. This may be the reason for which CHCl2 Br was below the detection limits during CCl3 Br degradation. The rate of disappearance are comparable for CHCl3 and CHBr3 both under air or helium, while d[X]/dt is twice −d[CX]/dt. For CHCl2 Br, the degra-

Fig. 7. (A) Photocatalytic degradation of 6.2×10−5 M CHCl2 Br on TiO2 200 mg l−1 at pH 5.5 in the presence of oxygen, the evolution of bromide and chloride ions and the formation of intermediates; (B) photocatalytic degradation of 6.2×10−5 M CHCl2 Br on TiO2 200 mg l−1 at pH 5.5 in the presence of helium, the production of bromide and chloride ions and the formation of intermediates.

P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34

dation rate is three times faster than CHCl3 and CHBr3 , while dX− /dt is close to d[CX]/dt. The rates of production of chloride and bromide ions under both experimental conditions are very similar. Taking into account the 2:1 Br:Cl ratio, bromide is released faster than chloride, while after 2 h of irradiation, the two halogens have still not reached the stoichiometric concentrations. Under helium, after 1 h, an almost complete disappearance of CHCl2 Br is obtained and Cl− and Br− are formed at 50 and 75% of the stoichiometric concentration. During longer irradiation times, halides are produced via a different kinetics, most likely due to the formation of halogenated intermediates. The rate of release of halide suggests that the degradation of CHCl2 Br passes through the formation, not only of the • CHCl2 radical, but also the • CHClBr radical due to the loss of chloride ions. This is supported by the formation of bromochloromethane. The • CHCl2 radical leads to the formation of CH2 Cl2 , while • CHClBr radical forms CH2 ClBr. The great amount of formic acid observed underlines that a sequence of oxidative and reductive attacks occurs. In aerated solution the rate of CHCl2 Br disappearance is very similar to the one followed in helium. The complete disappearance of the primary compound is observed after 2 h, while bromide ions reach almost the stoichiometric concentration (92%). Even though the rate of disappearance does not change, some notable differences are observed in structures and concentrations of the intermediates identified (Fig. 7b). In air-equilibrated system other steps involving oxygen are involved. In air, one of the halogenated intermediates identified is CH2 Cl2 . The absence of species containing bromine suggests that under these experimental conditions, the degradation of CHCl2 Br occurs preferentially through the formation of • CHCl2 radical. Referring to Scheme 1, carbon dioxide can be obtained via two separate oxidative pathways. In the presence of oxygen, the formation of • CCl2 Br radical is probably favoured, with respect to the • CHCl2 and • CHClBr radicals (the presence of oxygen reduces the free electrons), as is evidenced by the decrease of formaldehyde. CH2 Cl2 , CH2 ClBr and HCHO are formed through two reductive steps, while HCOOH is produced through combined oxidative–reductive paths.

31

CHCl2 radicals can undergo either reduction leading to the formation of CH2 Cl2 or oxidation resulting in :CCl2 formation. CH2 Cl2 can lead to HCHO, while :CCl2 produces formic acid. The formation of both species suggests that both steps are followed, while the large amount of formic acid (nC =2, as CHCl2 Br) implies that an oxido-reductive balanced sequence is realized easier than a purely reductive pathway. The presence of oxygen is compatible with a mechanism that favours the oxidative steps. The average of nC as a function of irradiation time is shown in Fig. 6. Under oxygen-free conditions, nC remains nearly constant, emphasizing the importance of both oxidative and reductive paths, similar to that seen for CHBr3 and CHCl3 . The presence of oxygen increases the likelihood of oxidative steps and nC increases toward +4. As mentioned above, the calculation of nC may be affected by increasing uncertainty, as the total amount of identified intermediates decreases. In these cases only the calculation at early times (30 min) are reported in Fig. 6. Anyway after 30 min 80% of the initial compound is disappeared. 3.3. Photocatalytic degradation of bromochloromethane The photocatalytic degradation of CH2 ClBr under aerobic and anaerobic conditions, the production of halide and the formation of intermediates is reported in Fig. 8a (air system) and in Fig. 8b (heliumsystem). In oxygen-free conditions, the rate of disappearance of CH2 ClBr is decreased and, after 4 h, only 50% of the primary compound disappears. These results are in agreement with those previously reported for CH2 Cl2 and CH2 Br2 . For these molecules, oxidative steps are favoured. For comparison, the degradation rates of CH2 Cl2 and CH2 Br2 are reported in Table 3. In all of these halogenated compounds the rate of disappearance decreases under anaerobic conditions. This is more pronounced for CH2 Cl2 , while for CH2 ClBr, the presence or absence of oxygen has a minor effect on the initial degradation rate. In the presence of helium, the initial degradation rate for the three halogenated compounds are similar, while in air CH2 Cl2 is degraded three times faster than the other two molecules.

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P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34 Table 3 Initial degradation rate of halomethanes (initial concentration CH2 Br2 5.7×10−5 M, CH2 Cl2 1.2×10−4 M, CH2 ClBr 7.8×10−5 M) on TiO2 200 mg l−1 under air or helium, and the corresponding initial rate of halide production Rate (M h−1 )×105 (19)

CH2 Cl2 Cl− CH2 Br2 (15) Br− CH2 ClBr Cl− Br−

Air

Helium

21.0 11.0 7.2 13.2 7.0 4.5 4.0

3.8 5.0 3.2 3.9 4.7 4.5 1.9

In the presence of air, chlorobromomethane can also react through reactions 21–23 CH2 ClBr + • OH → • CHClBr + H2 O

(21)

• CHClBr

+ • OH → HOCHClBr

(22)

• CHClBr

+ O2 → • O2 CHClBr

(23)

• CHClBr

radical can react with oxygen, leading to which breaks down to form CO2 , chloride and bromide. In the case of CH2 ClBr, both in the presence and in the absence of oxygen, formic acid and formaldehyde are formed at different concentrations. Under aerobic conditions, the concentrations of HCHO and HCOOH are formed in similar concentrations, while under anaerobic conditions, the main intermediate is formaldehyde. Such a difference can be understood by considering the mechanism of formation for these oxygenated compounds. Formic acid contains the carbon atom in +2 oxidation state and needs two oxidative attacks, while formaldehyde and CH2 ClBr contain the carbon atom in the same oxidation state (0) and need an oxidation–reduction sequence. In an oxygen-free system, the reductive/oxidative steps are balanced and formaldehyde is the principal intermediate, while the presence of oxygen enhances the oxidative steps and favours the formation of formic acid. Considering Fig. 6, nC appears to increase both under aerobic and anaerobic conditions, indicating the importance of oxidative pathways. Under a helium atmosphere, nC calculation is not reliable, because of the small percentage of intermediates identified. The increase in nC is particularly evident in air, where after 4 h nC =3. • O CHClBr, 2

Fig. 8. (A) Photocatalytic degradation of 7.8×10−5 M CH2 ClBr on TiO2 200 mg l−1 at pH 5.5 in the presence of oxygen, the production of bromide and chloride ions and the formation of intermediates; (B) photocatalytic degradation of 7.8×10−5 M CH2 ClBr on TiO2 200 mg l−1 at pH 5.5 in the presence of helium, the production of bromide and chloride ions and the formation of intermediates.

In Table 3, halide formation is also reported. For CH2 Cl2 , the rate of chloride production is higher in helium than in air, while for CH2 Br2 , the rate of bromide production is very high under both experimental conditions. Interestingly, under anaerobic conditions, chloride ions have been easily released (Cl:Br∼2:1) from CH2 ClBr, while in air the rate of chloride and bromide productions are very close. In the presence of air, 90% disappearance of CH2 ClBr was obtained after 4 h of irradiation. Whereas, a first-order kinetic is followed during the first hour of irradiation, in longer irradiation times a different kinetic is followed. Chloride and bromide were formed at similar rates and, in correspondence with complete CH2 ClBr degradation, stoichiometric concentrations of chloride and bromide ions were achieved.

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4. Conclusions The halomethanes studied have shown a complete disappearance of the primary compound and the halide formation in stoichiometric concentration, even if different rates are followed. Comparing the completely halogenated compounds, it emerges that under air CCl4 , CBr4 and CCl3 Br degrade with similar degradation rate, while the rate followed by CBr3 F is slower; in this case, the presence of fluorine stabilizes the molecule. The halides production is two–four times the rate of disappearance of the halomethanes. This may originate from the reaction of • CX3 radical with O2 and further fast halide release. Under helium the reductive steps are favoured and the degradation rates for the fully halogenated molecules increase dramatically with the exception of CCl3 Br, for which the rate is only twice. The asymmetry of the molecule decreases the rate; in fact, the degradation rate for CCl3 Br (11.7×10−5 M h−1 ) is slower than CCl4 (24×10−5 M h−1 ) and CBr4 ; the degradation rate of CBr3 F (7.3×10−5 M h−1 ) is slower than CBr4 (7.8×10−4 M h−1 ). The production rate of X− is comparable with halomethane rate of disappearance suggesting that the initial step involves the reaction with an electron bringing to X− release, while the other X− are released subsequently. The addition of methanol increases the degradation rate two–three times in all cases, while dX/dt is similar to d[CX3 Y]/dt, as previously seen under helium. Considering CHCl2 Br, the presence of a hydrogen atom increases the degradation rate two–three times, in respect to CCl3 Br, either in air and in helium. This is probably linked to the possibility for the molecule to degrade through oxidative and reductive paths, while for the fully halogenated molecules the initial degradation step involves necessarily a reductive attack. The asymmetry increases d[CHX2 Y]/dt. In fact, while for CHCl3 and CHBr3 the degradation rate are very similar (6.5 for CHCl3 , 7.2 for CHBr3 ) for CHCl2 Br is 20.6×10−5 M h−1 . Comparing the hydrohalomethanes either under air or helium, the degradation rates are quite similar. This is due to a kinetic balance between reductive and oxidative paths. Comparing the rate of CHCl3 with CCl4 and of CHBr3 with CBr4 , it emerges that an increase in 1.5 times in air is observed but a decrease is dramatically operating under helium. While the rate of halide production

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is comparable to d[CHX2 Cl]/dt it is about twice for CHCl3 and CHBr3 . A further substitution of hydrogen (CH2 ClBr) decreases the reactivity either under air and helium with respect to CHCl2 Br. Under helium the degradation rate decreases for the three dihydrohalomethanes considered underlyning that in this case the oxidative steps are predominant.

Acknowledgements The financial supports from CNR, MURST, Brite-Euram project #BE96-3593 Solartedox, Progetto Antartide and the Greece-Italy Exchange Program are kindly acknowledged. References [1] Scientific Assessment of Stratospheric Ozone, 1989, Global Ozone Research and Monitoring Project, Report No. 20, World Meteorological Organization, Vol. I, 1989. [2] W. Reifenhauser, K.G., Chemosphere 24 (1992) 1293. [3] R.A. Rasmussen, M.A.K. Khalil, Geophys. Res. Lett. 11 (1993) 433. [4] T. Class, R. Kohlne, K. Ballschmiter, Chemosphere 15 (1986) 429. [5] T. Class, K. Ballschmiter, J. Atmos. Chem. 6 (1988) 35. [6] M.B. Mc Elroy, R.J. Salawitch, S.C. Wofsy, J.A. Logan, Nature 321 (1986) 535. [7] N. Washida, T. Imamura, H. Bandow, Bull. Chem. Soc. Jpn. 69 (1996) 535. [8] T.M. Vogel, C.S. Criddle, P.L. Mc Carty, Environ. Sci. Tecnol. 21 (1987) 722. [9] W. Choi, M.R. Hoffmann, Environ. Sci. Technol. 31 (1997) 89. [10] M. Halmann, A.J. Hunt, D. Spath, Solar Energy Mater. Solar Cells 26 (1982) 1. [11] D. Mass, P. Pichat, C. Guillard, Res. Chem. Intermed. 23 (1997) 275. [12] D.W. Bahnemann, J. Monig, R. Chapman, J. Phys. Chem. 91 (1987) 3782. [13] P. Calza, C. Minero, E. Pelizzetti, Environ. Sci. Technol. 31 (1997) 2198. [14] E. Pelizzetti, C. Minero, V. Maurino, A. Sclafani, H. Hidaka, N. Serpone, Environ. Sci. Technol. 23 (1989) 1380. [15] C. Minero, E. Pelizzetti, S. Malato, J. Blanco, Chemosphere 26 (1993) 2103. [16] P. Piccinini, P. Calza, C. Minero, E. Pelizzetti, New J. Chem. 20 (1996) 1159. [17] P. Calza, C. Minero, A. Hiskia, E. Papaconstantinou, E. Pelizzetti, Appl. Catal. B: Environ. 21 (1999) 196. [18] P.J. Kropp, Acc. Chem. Res. 17 (1984) 131.

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P. Calza et al. / Applied Catalysis B: Environmental 29 (2001) 23–34

[19] C.E. Castro, N.O. Belser, J. Agric. Food Chem. 29 (1981) 1005. [20] C.E. Castro, S. Mayorya, N.O. Belser, J. Agric. Food Chem. 35 (1987) 865. [21] P. Calza, C. Minero, E. Pelizzetti, J. Chem. Soc., Faraday Trans. 93 (21) (1997) 3765. [22] G. Bullock, R. Cooper, Trans. Faraday Soc. 66 (1970) 2055. [23] P. Neta, R.W. Fessenden, R.H. Schuler, J. Phys. Chem. 75 (1971) 455, 1654. [24] J.F. Gibson, D.J.E. Ingram, M.C.R. Symons, M.G. Towsend, Trans. Faraday Soc. 53 (1957) 924. [25] V.J. Lilie, G. Beck, A. Henglein, Ber. Bunsenges. Phys. Chem. 75 (1971) 458.

[26] O. Micic, Y. Zhang, K.R. Cromack, A.D. Trifunac, M.C. Thurnauer, J. Phys. Chem. 97 (1993) 13284. [27] R. Mertens, C. Von Sonntag, J. Lind, G. Merenyi, Angew. Chem. Int. Ed. Eng. 33 (1994) 1259. [28] W. Choi, M.R. Hoffmann, J. Phys. Chem. 100 (1996) 2161. [29] K.D. Asmus, D.W. Bahnemann, K. Krisher, K. Lal, J. Honig, Life Chem. Rep. 5 (1985) 1. [30] N. Getoff, Water Res. 20 (1986) 1261. [31] N. Getoff, Radiat. Phys. Chem. 37 (1991) 673. [32] X. Shen, J. Lind, T.E. Eriksen, G. Merenyi, J. Phys. Chem. 93 (1989) 555.