Ind. Eng. Chem. Res. 1999, 38, 3345-3352
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Gas-Phase Reaction of Halon 1301 (CBrF3) with Methane Kai Li, Eric M. Kennedy,* and Bogdan Z. Dlugogorski Industrial Safety & Environment Protection Group, Department of Chemical Engineering, The University of Newcastle, Callaghan, NSW 2308, Australia
Nonoxidative gas-phase reaction of halon 1301 (CBrF3) with methane in a nitrogen bath was investigated using a tubular plug flow reactor. Experiments were performed at atmospheric pressure, over a range of temperatures (673-1053 K) and residence times (0.1-2.0 s). Compared to the thermal decomposition of CBrF3, the addition of CH4 to the reacting stream results in a substantial increase in the conversion of CBrF3, with conversion levels rising with increasing content of CH4. Generally, the conversion of both reactants increases with temperature or residence time. At high temperatures and an equal-molar CBrF3/CH4 feed stream, the proportion of CBrF3 converted is always greater than that of CH4. In addition to HBr and HF, the major products of the reaction were CHF3, CH3Br, and C2H2F2, while minor products include C2H4, C2H2, C2H3Br, CHBrF2, C2F6, C2H3F3, C2HBrF2, C2H3F, C2HF5, C6H5F, C6H5BrF, CH2Br2, and H2. Coke formation was observed above 960 K. A reaction mechanism for prediction of major and important minor species is presented and discussed. The reaction kinetics can be represented (at low conversion) by a second-order global reaction scheme with the following rate parameters: kglobal ) 3.41 × 1015 (cm3‚mol-1‚s-1) exp(-180.2 (kJ‚mol-1)/RT). Introduction Halon fire suppression agents, such as halons 1301 (CBrF3), 1211 (CBrClF2), and 2402 (C2Br2F4), are bromine- or chlorine-containing fluorocarbons. They have been used extensively for fire mitigation applications. However halons, like CFCs, have very long tropospheric lifetimes and, consequently, can diffuse into the stratosphere. In the stratosphere, halons and CFCs are broken down by UV light, resulting in the formation of Cl and Br atoms, which catalyze the depletion of the ozone layer, the earth’s primary shield against harmful UV radiation (Molina and Rowland, 1974). In fact, Br atoms are suggested to be 10-100 times as destructive as chlorine atoms on a per-atom basis (McElroy and Salawich, 1989). The Montreal protocol has mandated a ban on the production of halons and other ozone depletion substances in industrialized countries. In addition, Australia has not only phased out the production of halons but also banned their use. Australia has replaced CFCs and halons in all but a small number of specific applications and has introduced replacement compounds whose ODP (ozone depletion potential) are very low or essentially zero. The removal of halons from use has resulted in two problems: one related to the need to treat the growing stockpiles of halons and the second related to the need of developing suitable replacements for these compounds. The current focus of worldwide research is on the development of treatment technologies for halons, where destruction and conversion are parallel alternative treatment approaches. Destruction technologies consist of thermal, chemical, and electrical processes. With the exception of thermal processes, other technologies that could (potentially) be adopted to destroy halons are still at the bench-scale * To whom correspondence should be addressed. Phone: +61 2 4921-6177. Fax: +61 2 4921-6920. E-mail: cgek@ cc.newcastle.edu.au.
level of development (Dickerman et al., 1989). At present, the only established treatment technology is based on incineration, including rotary kiln, fluidizedbed, liquid injection, lime/cement kiln coincineration, molten salt incineration, plasma arc, and high-temperature fluidized bed destruction (Turner, 1986). These thermal destruction processes expose halons to high temperatures at long residence times in the presence of excess oxygen. However, the combustion inhibition properties of halons make incineration a very unattractive disposal option from an energy consumption perspective. Moreover, at the high temperature of incineration, halons are broken down into corrosive halogen acids (HF, HCl, HBr) and free halogen molecules (F2, Cl2, Br2) which pose significant sample handling problems. In particular, hydrogen fluoride (HF) aggressively attacks the refractory of many current incinerator systems (Dickerman et al., 1989). Furthermore, there are increasing concerns over the emission of both feed and the carcinogenic or toxic products of incomplete combustion (PICs), which accompany many oxidative processes. There are a limited number of studies of catalyst-aided oxidative destruction of halons at bench scale, but catalyst deactivation remains an unresolved problem in the development of this technology (Bickle et al., 1994; Karmakar and Greene, 1994; Nagata et al., 1994; Tajima et al., 1997). Our research interest has focused on investigating hydrodehalogenation as a possible technology for treatment of halons, where the focus is on conversion of halons into compounds of economic value, such as HFCs, which have been introduced as substitutes for halons. Nonoxidative thermal hydrodehalogenation (THD), where halons or halogenated compounds react with hydrogen or hydrogen donors at high temperature, is recognized as an emerging treatment process. De Lijser et al. (1994) have investigated THD of bromochlorodifluoromethane (CBrClF2, halon 1211) in hydrogen (H2) and have found that, at low temperature (95%) of CBrF3 occurs at about 1023 K and at a residence time of 2.0 s, while the corresponding conversion of CH4 was approximately 72%. Comparison of the data in Figures 3 and 4 clearly shows that, at low temperatures, there is little difference between conversion levels of CBrF3 and CH4, while at high temperatures, the conversion level of CBrF3 is greater than that of CH4.
Figure 5. CBrF3 conversion as a function of temperature at 1.0 s residence time at various CF3Br/CH4 ratios.
Two possible explanations for the lower conversion of CH4 compared with CBrF3 can be presented. First, less CH4 is required to react with CBrF3 at high temperatures, because more than one H-C bonds of CH4 can be activated, and thus one molecule of CH4 can contribute more than one hydrogen atom to per molecule of CBrF3. Second, reaction products, such as HBr and C2H4, whose concentrations increase with temperature, compete with CH4 to react with CBrF3, resulting in a decrease in the relative amount of CH4 consumed. HBr, in particular, probably does not play a significant role at low temperature in the reaction with CBrF3 due to its low concentration, but its role becomes significant at higher temperature. HBr has been identified as a very efficient hydrogen transfer agent (Russel et al., 1988). For example, the reaction of HBr with CF3 is about 25 times faster than the corresponding reaction of CF3 with H2 at 800 K (Weeks and Whittle, 1983). It can be seen from Figure 5 that, at a 1.0 s residence time and in the absence of CH4, the conversion level of CBrF3 was approximately 10% at 1053 K. The addition of a small amount of methane dramatically increases the conversion of CBrF3. Moreover, the conversion of CBrF3 increases with increasing concentration of CH4 in the feed, and this effect is more pronounced at high temperatures. Clearly, the presence of CH4 facilitates the decomposition of CBrF3. A similar effect has been observed by Ritter (1994), who studied CF2ClCFCl2 decomposition in the presence of hydrogen. Conversely, Graham and others (1986) investigated the reaction of CFC 113 under inert and oxidative conditions, but they found no difference in reaction activity. Pitts et al. (1990) and Ritter (1994) compared oxidation reactions with hydrogenation reactions for halogenated hydrocarbons and concluded that perhalogenated hydrocarbons are inert to attack by O or OH radicals but are favorable to attack by H radicals. This hypothesis was further confirmed by our results, and we suggest that perhalogenated compounds are also favorable to attack by CH3 radicals. The data presented suggest that, at low temperature, the thermal cleavage of C-Br bond, the most labile bond in CBrF3 (Babushok et al., 1996), initiates a radical chain reaction, where the CF3 radicals attack CH4 to produce CHF3 and a CH3 radical. Once the CH3 radical is formed, it catalyzes the conversion of CF3Br into CF3 via reaction R3, forming another major product species,
3348 Ind. Eng. Chem. Res., Vol. 38, No. 9, 1999
Figure 6. Proposed cycle for the reaction of CBrF3 with CH4 at low temperatures.
Figure 8. Arrhenius plot for 2nd order reaction rate constant for the coupling reaction of CBrF3 and CH4.
Figure 7. Apparent 2nd order behavior for the coupling reaction of CBrF3 and CH4 over the temperature range 723-903 K. X represents fractional conversion of CBrF3.
CH3Br. At low temperature, the reaction cycle would be as shown pictorially in Figure 6.
CF3Br f CF3 + Br
(R1)
CF3 + CH4 f CHF3 + CH3
(R2)
CH3 + CF3Br f CH3Br + CF3
(R3)
It should be pointed out that the reaction of CBrF3 with CH4 is initiated by reaction R1, rather than by the decomposition of CH4. However, the key pathway for production of the CF3 radical is via reaction R3 and not reaction R1. Thus, without CH4 added in the feed stream, the decomposition of CBrF3 is very slow. With addition of CH4, reactions R2 and R3 are fast and promote a rapid conversion of CBrF3 to CH3Br and CHF3. Global Reaction Kinetics. From the conversion data obtained, global reaction kinetics was estimated. Equimolar concentrations of CBrF3 and CH4 were used in the feed stream, diluted by N2 (the CH4-rich and CBrF3-rich cases were not considered here). Hence, at low conversion, the reaction rate could be expressed as
-rA ) kCaACbB ) kCnA n ) a + b (global reaction order) (3)
Figure 9. GCMS chromatograph trace of effluent species from the reaction of CBrF3 and CH4 at 993 K. Key: 1, N2 and CH4; 2, C2F6 and C2HF5; 3, CHF3; 4, C2H4; 5, C2H2; 6, C2H2F2; 7, C2H3F; 8, C2H3F3; 9, CBrF3; 10, C6H5F; 11, CHBrF2; 12, CH2BrF; 13, CH3Br; 14, C2HBrF2; 15, C2H3Br; 16, CH2Br2; 17, C6H4BrF.
where k is the reaction rate constant ((mol/cm3)1-n/s), CA is the concentration of CF3Br (mol/cm3), CB is the concentration of CH4 (mol/cm3), and -rA is the rate of the reaction (mol/cm3‚s). We evaluated the global kinetics of the reaction, using the ideal plug-flow equation and assuming a constantdensity system for simplicity. The first-order and secondorder equations were used to fit the experimental data, and the results showed that the global reaction between CBrF3 and CH4 follows second-order kinetics throughout
Ind. Eng. Chem. Res., Vol. 38, No. 9, 1999 3349
Figure 10. Major product distribution as a function of reaction temperature at a 1.0 s residence time, with product yield expressed as fractions of the initial mole fraction of CBrF3. Volumetric ratio of reactant feed: N2:CBrF3:CH4 ) 11:1:1.
the temperature range studied. Arrhenius plots of X/(1 - X) versus residence time τ are shown in Figure 7. On the basis of the data extracted from Figure 7, an Arrhenius plot of rate constant vs 1/T (Figure 8) was produced, and the second-order rate constant parameters obtained were
In equation R4, M is collision partner. Biordi et al. (1978) found that the decomposition reaction of CHF3 is relatively slow, when they investigated the CBrF3inhibited methane flames. They indicated that reaction R5 is of primary importance in the high-temperature destruction of CHF3. As for the consumption of CH3Br, reaction R9 is believed to be a very minor contribution, based on calculations by Biordi et al. (1978). Reaction R10 is relatively unimportant as well, as it is too slow compared with its reverse reaction. For example, at 1000 K, the rate constant of reaction R10 is k10 ) 1.70 × 105 cm3 mol-1 s-1, while its reverse reaction rate is k′10 ) 9.97 × 1012 cm3 mol-1 s-1 (Westbrook, 1983). The primary mechanism for the consumption of CH3Br is likely to proceed via reactions R7 and R8, which is in agreement with the observation by Battin-Leclerc et al. (1994), who found that 72% CH3Br was consumed by (R8) and 24% CH3Br by (R7) at 1070 K and 0.5 s when they investigated the inhibiting effect of CBrF3 on methane flames. The yield of other species, including C2H2F2, increases with temperature over the entire temperature range. It is generally agreed that C2H2F2 is formed from radicals CH3 and CF3, but there is dispute concerning the specific pathways to C2H2F2 formation. Westbrook (Westbrook, 1983) reported that C2H2F2 is formed directly from CF3 + CH3, through elimination of HF.
CF3 + CH3 f CH2dCF2 + HF
kglobal ) 3.41 × 1015 (cm3‚mol-1‚s-1) × exp(-180.2 (kJ‚mol-1)/RT) (4) Product Variation with Temperature. In addition to HBr and HF, the major products formed during reaction were CHF3, CH3Br, and C2H2F2. Minor products species included C2H4, C2H2, C2H3F3, C2HBrF2, C2H3F, C2HF5, C2F6, C6H5F, CHBrF2, and C2H3Br (Figure 9). Soot and trace quantities of C6H4BrF, Br2, and H2 were also detected at high temperatures. Only two products, CHF3 and CH3Br, were detected at temperatures up to 873 K. The variation of major products with temperature at a constant residence time of 1.0 s is shown in Figure 10. It is evident that the yield of CHF3 and CH3Br increases with temperature but reaches a maximum at a temperature of 970 K for CH3Br and 1020 K for CHF3. The maximum temperature for CH3Br is lower than that of CHF3, consistent with the lower dissociation energy of C-Br (293 ( 5 kJ/mol) compared to C-H (444 ( 4 kJ/mol) or C-F (452 ( 13 kJ/mol) bonds (Weast, 1985). Both CHF3 and CH3Br can decompose or react further at higher temperature, leading to consumption of these initial products. The following reaction scheme for decomposition of CHF3 and CH3Br has been suggested (Biordi, et al., 1978, Westbrook, 1983):
(R11)
Biordi et al. (1978) and Battin-Leclerc et al. (1994), on the other hand, consider two reaction steps leading to the formation of C2H2F2.
CF3 + CH3 f CF3CH3
(R12)
CF3CH3 f CH2CF2 + HF
(R13)
Since CF3CH3 was detected (albeit in small quantities) experimentally in this study, the latter two step mechanism for C2H2F2 formation seems to be favorable under these reaction conditions. The reaction
CF2 + CH3 f CH2CF2 + H
(R14)
seems to be a plausible mechanism to contribute to the formation of C2H2F2, via vibrationally excited CH3CF2, but Biordi et al. (1978) suggest that its contribution toward formation of C2H2F2 is minimal. HBr was detected in small amounts at low temperature, but its concentration increases rapidly at high temperatures. At low temperature, the Br radicals and Br2 molecules are consumed in reactions R15-R21. In
CH3 + Br f CH3Br
(R15)
CHF3 + M f CF3 + H + M
(R4)
Br + CH4 f CH3 + HBr
(R16)
CHF3 f CF2 + HF
(R5)
CH3 + HBr f CH4 + Br
(R17)
CHF3 + H f HF + CHF2
(R6)
Br + Br + M f Br2 + M
(R18)
CH3Br f CH3 + Br
(R7)
CH3 + Br2 f CH3Br + Br
(R19)
CH3Br + H f HBr + CH3
(R8)
CF3 + Br2 f CBrF3 + Br
(R20)
CH3Br + Br f HBr + CH2Br
(R9)
Br + H + Mf HBr + M
(R21)
CH3Br + Br f Br2 + CH3
(R10)
these reactions, only (R16) forms HBr, but this reaction
3350 Ind. Eng. Chem. Res., Vol. 38, No. 9, 1999
is very slow at low temperatures since the reverse reaction R17 is very fast due to its negative activation energy (Russel et al., 1988). Conversely, reaction R16 is greatly enhanced at high temperature, and even reactions R8, R9, and R21 promote HBr formation at high temperature. Reaction R8, in particular, is very fast at high temperatures compared with other CH3Br consumption reactions and is one of the major hightemperature channels for production of HBr. HF formation was detected above 933 K, and its concentration increased substantially with temperature, which agrees well with previous studies (de Lijser et al., 1994; Ritter, 1994). In addition to reactions R5, R6, and R13, HF can also be generated in reactions R22R24.
CF3 + H f CF2 + HF
(R22)
CF2 + H f CF + HF
(R23)
CF + H f C + HF
(R24)
Figure 11. Minor product distribution as a function of reaction temperature at 1.0 s residence time, with product yield expressed as fractions of the initial mole fraction of CBrF3. Volumetric ratio of reactant feed: N2:CBrF3:CH4 ) 11:1:1.
Among these reactions, (R22) is believed to be the most important reaction for HF generation, due to the high concentration of CF3 radicals in the reaction pool. Reactions R22-R24 are all energetically favorable; their reaction heats (∆H) are -202, -53, and -29 kJ/mol, respectively (Stull and Prophet, 1970). However, the effective rates of (R23) and (R24) are lower than that of (R22) because of low concentration levels of CF2 and CF radicals. The reaction chemistry of (R22)-(R24) has been the focus of intensive research interest for many years (Tsai and McFadden, 1989; Biordi et al., 1978). It has been suggested that there is a low-energy potential surface, on which those reactions occur through atom-radical addition to form a stable intermediate complex such as CHF3 or CHF2, and then the subsequent breakup of these complexes results in the formation of HF (Tsai and McFadden, 1989). It is also possible that the fluorine atoms are abstracted directly by hydrogen atoms, leading to the production of HF (Tsai and McFadden, 1989). Conversely, it has been suggested that these reactions may result directly in F atoms, which rapidly abstract hydrogen atoms from other species to yield HF (Biordi et al., 1978). Reactions R5 and R13 are the major contributors to production of HF, because (R5) is a major CHF3 consumption reaction and (R13) is a major contributor to formation of C2H2F2. There are other possible reaction steps such as (R25)(R29), which also lead to the production of HF, but it is believed that they only play a very minor role, either because reacting species are present in very low concentration or they are energetically unfavorable.
species but were only present in low concentrations (
