Oxidation of Dimethyl Ether: Absolute Rate Constants for the Self Reaction of CH 3. OCH 2 Radicals, the Reaction of CH 3.

Transcription

1 Oxidation of Dimethyl Ether: Absolute Rate Constants for the Self Reaction of Radicals, the Reaction of Radicals with, and the Thermal Decomposition of Radicals JENS SEHESTED, KNUD SEHESTED, JESPER PLATZ, HELGE EGSGAARD, OLE JOHN NIELSEN Section for Chemical Reactivity, Environmental Science and Technology Department, Risø National Laboratory, DK-4000 Roskilde, Denmark Received 10 October 1996; accepted 6 January 1997 ABSTRACT: The rate constant for the reaction of radicals with (reaction (1)) and the self reaction of radicals (reaction (5)) were measured using pulse radiolysis coupled with time resolved UV absorption spectroscopy. k 1 was studied at 296 K over the pressure range bar and in the temperature range K at 18 bar total pressure. Reaction (1) is known to proceed through the following mechanism: ;: # 9: CH 2 H # 9: 2HCHO OH ;: # M 9: M (k prod ) (k RO2 ) k 1 k RO2 k prod, where k RO2 is the rate constant for peroxy radical production and k prod is the rate constant for formaldehyde production. The k 1 values obtained at 296 K together with the available literature values for k 1 determined at low pressures were fitted using a modified Lindemann mechanism and the following parameters were obtained: k cm 6 molecule 2 s 1 k cm 3 RO2, ( ) RO2,0 ( ) 10 30, molecule 1 s 1, and k cm 3 molecule 1 s 1 prod,0 ( ) where k RO2,0 and k RO2, are the overall termolecular and bimolecular rate constants for formation of rad-

2 628 SEHESTED ET AL. icals and k prod,0 represents the bimolecular rate constant for the reaction of radicals with to yield formaldehyde in the limit of low pressure. k RO2, ( ) exp( (46 27)/ T) cm 3 molecule 1 s 1 was determined at 18 bar total pressure over the temperature range K. At 1 bar total pressure and 296 K, k cm 3 5 ( ) molecule 1 s 1 and at 18 bar total pressure over the temperature range K, k cm 3 molecule 1 s 1 5 ( ) As a part of this study the decay rate of radicals was used to study the thermal decomposition of radicals in the temperature range K at 18 bar total pressure. The observed decay rates of radicals were consistent with the literature value of k exp[ 12800/ T) s 1. The results are discussed in the context of dimethyl ether as an alternative diesel fuel John Wiley & Sons, Inc. Int J Chem Kinet 29: , INTRODUCTION # M 9: M (1c) Topsøe and Amoco have recently proposed dimethyl ether (DME), O, as a diesel substitute [1]. Dimethyl ether has several properties which make it an attractive diesel fuel. Firstly, it has a low selfignition temperature and a high cetane number (235 C, 55 60). Secondly, when used in diesel engines it reduces combustion noise and NO x emissions. Thirdly, DME fueled engines are essentially nonsooting. Fourthly, DME can be produced from C 1 feedstocks via a low cost one step synthesis [1]. Engine tests have shown that DME fueled diesel engines have emission levels that surpass the California Ultra Low Emissions Vehicle (ULEV) regulation for medium duty vehicles [2]. At this point it is germane to note that unlike conventional diesel fuel, DME is a gas at ambient temperature and pressure. Hence, any widespread use of DME would require significant investment in new infrastructure associated with fuel distribution and delivery. The atmospheric chemistry of DME has been studied by Japar et al. [3], Jenkin et al. [4], Wallington et al. [5], Langer et al. [6], and Sehested et al. [7]. Jenkin et al. [4] and Sehested et al. [7] reported high formaldehyde yields following chlorine initiated oxidation of dimethyl ether in a N 2 / diluent at total pressures below 300 torr. It was shown that formaldehyde is formed from the reaction of radicals with via the following mechanism [4,7]: ;: # # 9: CH 2 H # 9: 2HCHO OH (1a, 1a) (1b) Correspondence to: O. J. Nielsen Contract grant Sponsor: Danish Technical Research Council Contract grant Sponsor: Ford Motor Company 1997 John Wiley & Sons, Inc. CCC /97/ where formaldehyde is formed by intramolecular hydrogen transfer in the energetically excited # radical followed by decomposition, reaction (1b). At sufficiently low pressures ( 1 mbar) # radicals are not stabilized and reaction (1) proceeds essentially via reaction (1b). At higher pressures the highly energetic # complex is stabilized by collisional quenching to give thermalized radicals. radicals react to give OCHO and OOH under atmospheric conditions [3 5]. Alkylperoxy radicals are thermally unstable at 1000 K and consequently are not important in high temperature combustion chemistry. However, at lower temperature ( K) alkylperoxy radicals have sufficient stability to play an important role in unwanted autoignition phenomena in gasoline fueled engines ( knocking ) [8]. The importance of reaction (1) and (2) during the combustion of DME has been considered by Dagaut et al. [9]. ( M) 9: Products M 9: HCHO M (1) (2) Dagaut et al. [9] used combustion of DME in a jet stirred reactor at temperatures of K and pressures of 1 10 bar followed by detailed kinetic modeling of the obtained experimental data. It was found that reaction (1) was an important loss of radicals in their system at high pressures and low temperatures [9], e.g., under ignition conditions while reaction (2) was the dominant fate of radicals at combustion temperatures. In particular, reactions (1b) and (2) are important since they propagate the reaction, producing OH radicals

3 OXIDATION OF DIMETHYL ETHER 629 and radicals, respectively, and formaldehyde. Combustion of formaldehyde gives little, or no, particulate matter [10]. The primary objective of this work was to investigate the kinetics of reaction (1) and (2). A secondary objective was to study the kinetics of the self reaction of the radical. The results reported herein are discussed in the context of DME as an alternative diesel fuel. EXPERIMENTAL Two facilities using pulse radiolysis coupled with time resolved UV absorption spectroscopy were used. Both have been described in detail elsewhere [11,12] and will only be described briefly here. The 0.33 l cylindrical high pressure cell has a quartz window in one end of the cell and a 2 mm stainless steel window at the other end. Reactions were initiated by a 10 MeV HRC Linarc electron accelerator delivering a s electron pulse into the reaction cell. Partial pressures in the cell of less than 1 bar were measured by a MKS 122AA-01000AB pressure transducer coupled to a MKS PDR-C-2C instrument. Pressures above 1 bar were measured by a Juno 4AP-30 pressure transducer coupled to a 4PDE- 48 instrument. The cell was operated at a total pressure of 18 bar and at temperatures between 296 and 666 K. A pulsed 150 W Varian Xenon arc lamp delivered the analyzing light. The analyzing light entered the cell through a quartz window, and was reflected by a mirror mounted at the end of the cell, giving a total optical path length of 20 cm. The light was analyzed by a Perkin-Elmer double-quartz prism monochromator with an optical resolution of 2 5 nm, an IP28 photomultiplier, and a LeCroy 9400 digital oscilloscope. Data handling and storage were performed using a IBM personal computer. To improve S/N ratios, up to five pulses were averaged. No changes in the observed transients were detected by comparing the transients obtained from the first and last electron pulse. In the second setup, radicals were produced in a 1 L stainless steel reaction cell using a 30 ns pulse of 2 MeV electrons from a Febetron 705 B field emission accelerator. A total pressure range of bar and a temperature of 296 K was used in this setup. The pressure was measured by a Baratron absolute membrane manometer. The gas temperature was measured by a chromel/alumel thermocouple close to the center of the reaction cell. Reactions were monitored using the UV light from a pulsed 150 W Varian Xenon arc lamp. The light beam from the Xenon lamp was reflected in the gas cell by internal White-type optics, giving a total optical path length of 120 cm. The analyzing light was diffracted by a 1 m McPherson grating UV-VIS monochromator and detected with a Hamamatsu R928 photomultiplier. The monochromator was operated at a spectra resolution of 0.8 nm. Data handling was performed using a IBM personal computer. The gas mixtures used for both setups contained SF 6 in great excess. F atoms are known to be produced upon radiolysis of SF 6 with high-energy electrons [13]: 2 or 10 MeV e SF : F products (3) The F atom yield was determined in the low pressure system by radiolysis of gas mixtures of CH 4 / /SF 6 and the subsequent observation of the absorbance at 260 nm attributed to radicals. Using ( ) cm 2 molecule 1 [14] and the absorption at 260 nm, the F atom yield was determined to be ( ) cm 3. The F atom yield in the high pressure system was determined by radiolysis of mixtures of 100 mbar of DME and 18 bar of SF 6 to be (5 1) cm 3 for a 2 s electron pulse using the known absorption cross sections of radicals in the range nm [6]. Reagents used were: ultra high purity ( %), 0 6 mbar, supplied by L Air Liquide; SF 6 (99.97% and 99.9% using 18 and 1 bar total pressure, respectively), 20 mbar-18 bar, supplied by Messer Griesheim; and O ( 99.99%), mbar, obtained from AGA Gas AB. All reagents were used as received. Five sets of experiments were performed as a part of the present work: Firstly, using the low pressure setup, the rate constant for the reaction of radicals with, reaction (1), was measured in the pressure range bar total pressure (SF 6 ) at 296 K. Secondly, the rate constant for reaction (1) was determined at 18 bar total pressure (SF 6 ) in the temperature range K. Thirdly, the self reaction of the radicals was investigated at 1 bar total pressure (SF 6 ) and at 296 K. Fourthly, the rate constant for the self reaction of the radical was studied at 18 bar total pressure (SF 6 ) in the temperature range K. Finally, the thermally induced decomposition of the radical was investigated at 18 bar total pressure (SF 6 ) at temperatures between 573 to 666 K.

4 630 SEHESTED ET AL. RESULTS Rate Constants for the Reaction The reaction of radicals with was studied at 296 K and mbar total pressure using the low pressure setup and at K and 18 bar total pressure (SF 6 ) using the high pressure setup. Following pulse radiolysis of gas mixtures of 1 5 mbar of O, mbar of, and SF 6 added to 25, 50, 100, 200, and 1000 mbar total pressure a transient absorption was observed at 310 nm. Figure 1(A) and (B) show experimental transients obtained with a detection wavelength of 310 nm with 1000 and 50 mbar total pressure. The following reactions are expected to be important in the reaction system: 2 or 10 MeV e SF : F products F O 9: HF M 9: products M 9: products (3) (4) (5) (1) We attribute the transient absorption at 310 nm to formation and decay of radicals for two reasons. Firstly, the maximum transient absorption at 310 nm measured in this work is consistent with Langer et al. [6]. Secondly, the formation and decay of the transient absorption can be rationalized as follows from reactions (1,3 5). The fast rise is due to the formation of radicals from the reaction of F atoms with dimethyl ether, reaction (4). The decay was second order without present in the system due to reaction (5) and with present the decay rate increased with increasing [ ] due reaction (1). The absorption cross section of radicals is small at 310 nm [6] and the absorption transient therefore tends to essentially zero absorption. It seems reasonable to ascribe the transient absorption at 310 nm to the formation and decay of radicals. The rate constant for reaction (1) was determined by fitting the decay of the absorption by a first-order decay expression: Abs(t) (A 0 A inf )exp( k 1st t) A inf (I) Figure 1 Figure 1 show experimental transients obtained with a detection wavelength of 310 nm following radiolysis of 5 mbar DME, 3.78 mbar, and mbar of SF 6 (A) and 2 mbar DME, 1.22 mbar, and 46.8 mbar of SF 6 (B). The optical path length was 120 cm. The smooth solid lines are first-order fits to the experimental data and the dashed line in (B) is a simulation of the absorption transient, see text for details. where Abs(t) is the time dependent absorbance at 310 nm, and the three parameters A 0, A inf, and k 1st are the absorbance of radicals extrapolated back to time zero, the absorbance extrapolated to infinite time, and the pseudo-first-order decay rate for the radical, respectively. Expression (I) gives a good fit to the experimental data as seen from Figure 1. The values of k 1st obtained are plotted in Figure 2(A) as a function of the initial concentration of. As shown in Figure 2(A) the values of k 1st are proportional to [ ] for each total pressure. The y-axis intercepts are due to the self reaction of radicals, reaction (5). The slope of the straight lines through the data give the rate constant for reaction (1) at each total pressure. The values obtained are k 1 ( ), ( ), ), ( ), and

5 OXIDATION OF DIMETHYL ETHER 631 al. [15]. The values of k 1 obtained by Masaki et al. [15] are independent of total pressure in the pressure range mbar. This is, however, in line with the results reported by Sehested et al. [7] who found that k 1 is independent of total pressure in the pressure range mbar within 25%. This was explained by the following mechanism for reaction (1): ;: # # 9: CH 2 H # 9: 2HCHO OH # M 9: M (1a, 1a) (1b) (1c) Figure 2 (A) Plot of the pseudo-first-order decay rate constants obtained at 296 K from fits to experimental absorption transients using 1000 (filled circles), 200 (hollow circles), 100 (filled squares), 50 (filled triangles), and 25 mbar (hollow triangles) total pressure plotted as function of the oxygen partial pressure. (B) Rate constants for the reaction of with plotted as function of total pressure. Rate constants obtained in this work are shown with filled circles while rate constants reported by Masaki et al. [15] and Hoyermann et al. [17] are plotted with hollow circles and a hollow triangle, respectively. The S- shaped smooth line is a fit to the experimental data using a modified Lindemann mechanism. See text for details. ( ) cm 3 molecule 1 s 1 at 1000, 200, 100, 50, and 25 mbar total pressure, respectively. These values are plotted as a function of total pressure in Figure 2(B). Also plotted in Figure 2(B) is k 1 determined at 18 bar total pressure using the high pressure setup as discussed below. As seen from Figure 2(B) reaction (1) is pressure dependent as expected for a third-order reaction. Also plotted in Figure 2(B) are the rate constants obtained by Masaki et At high pressures the rate constant for the reaction of radicals with equals k 1a since all O # 2 radicals are quenched by reaction (1c), and the peroxy radical,, is formed. As the total pressure is decreased not all of the O # 2 radicals are quenched and the rate constant k 1 decreases. At the low pressure limit none of the O # 2 radicals are quenched, however, because of the intramolecular hydrogen transfer in a six membered transition state of the excited O # 2 radical the overall rate constant k 1 does not drop to zero at zero total pressure. Reactions (1a 1c) can be viewed as a modified Lindemann mechanism. The overall reaction of radicals with proceeds via two pathways to give either the peroxy radical (R ) or other products (prod). Hence, d[ ]/dt d[r ]/dt d[prod]/dt k 1 [ ][ ] d[r ]/dt k RO2 [ ][ ] k 1c [M][R# ] d[prod]/dt k prod [ ][ ] k 1b [R# ] Applying a steady-state analysis for R # gives k RO2 k 1a k 1c [M]/(k 1a k 1b k 1c [M]) k prod k 1a k 1b /(k 1a k 1b k 1c [M]) Defining k RO2,0 k 1a k 1c /(k 1a k 1b ) and k RO2, k 1a we can express the effective second-order rate constant for the formation of R radicals at any given pressure, k RO2, in terms of the limiting low and high pressure rate constants, k RO2,0 and k RO2, : k RO2 (k 1a k 1c [M]/(k 1a k 1b ))/ (1 k 1a k 1c [M]/(k 1a (k 1a k 1b ))) k RO2,0 [M]/(1 k RO2,0 [M]/k RO2, )

6 632 SEHESTED ET AL. Similarly, defining the rate constant for formation of products (formaldehyde and OH radicals) in the low pressure limit, k prod,0 k 1a k 1b /(k 1a k 1b ) it follows that k prod (k 1a k 1b /(k 1a k 1b ))/(1 k 1a k 1c [M]/ (k 1a (k 1a k 1b ))) k prod,0 /(1 k RO2,0 [M]/k RO2, ) Expression (II) k 1 k RO2 k prod (II) was fitted to the values of k 1 obtained here and the values reported by Masaki et al. [15] plotted in Figure 2(B) with k prod,0, k RO2,0, and k RO2, as parameters and weighting the data with the reciprocal of the uncertainty. The fit is shown as the solid line in Figure 2(B). As seen from Figure 2 the experimental data are well fitted by eq. (II). The best fit was obtained using k cm 3 prod,0 ( ) molecule 1 s 1, k cm 6 RO2,0 ( ) molecule s 1, and k cm 3 RO2, ( ) molecule 1 s 1. As a check of these kinetic parameters, a model consisting of k 1 k RO2 k prod (this work), k (estimated), k (this work), k CH3OCH2O2 CH3OCH2O [4], k cm 3 molecule 1 s 1 OH CH3OCH [16], and 310nm ( CH cm 2 3 ) molecule [6] was used to simulate the absorption tran- 1 sient in Figure 1(B). As seen from Figure 1(B) there is excellent agreement between the simulated and the experimental data. This finding serves to validate our experimental technique. Sehested et al. [7] obtained values for k prod,0, and k RO2,0, and k RO2, relative to k CH3OCH2 C12 of k cm 3 RO2,0 /k CH3OCH2 C12 ( ) molecule 1, k RO2, /k CH3OCH2 C , and k prod,0 /k CH3OCH2 C12 ( ) Sehested et al. [7] determined k CH3OCH2 C12 using the absolute rate constants obtained in this work and those obtained by Masaki et al. [15]. Hence, only the rate constant ratios k RO2, /k prod,0 and k RO2, /k prod,0 obtained in this work and the work of Sehested et al. [7] can be compared directly. The values of k RO2, /k prod, from this work and k RO2, /k prod, obtained by Sehested et al. [7] are in good agreement. However, the value of k cm 3 RO2,0 /k prod,0 ( ) molecule s 1 obtained here is significantly lower than 1 the value of k cm 3 RO2,0 /k prod,0 ( ) molecule 1 s 1 reported by Sehested et al. [7]. It is expected that the value of k RO2,0 /k prod,0 determined here could be higher than that of Sehested et al. [7] due to the higher third body efficiency of SF 6 compared to N 2. However, it is the rate constant ratio k RO2,0 /k prod,0 determined using SF 6 as bathgas which is the lowest. The reason for this apparent discrepancy is unknown at present. Hoyermann et al. [17] determined k 1 at 1 5 mbar total pressure and in the temperature range K to be k 1 ( ) exp(( )/ T) cm 3 molecule 1 s 1. The value of k 1 obtained from this expression at 296 K is shown in Figure 2(B) for comparison. As seen from the figure, the value determined by Hoyermann et al. [17] is 60% higher than the values determined by Masaki et al. [15]. The reason for this discrepancy is unknown. Reaction (1) was also studied as a function of temperature in the temperature range K at 18 bar total pressure. The decay of the radical was monitored at 310 nm as the detection wavelength following addition of 0 6 mbar of. The decay was second-order when no was present. The kinetics of the decay changed to first-order when was added. The decay of the transient absorptions was fitted with a first-order expression, expression (I), and pseudo-first-order rate constants, k 1st, were obtained. In Figure 3(A) the pseudo-first-order rate constants are plotted as a function of the oxygen concentration. k 1 was measured at four different temperatures: 296, 353, 423, and 473 K. We obtained the following rate constants: k 1 (296 K) ( ), k 1 (353 K) ( ), k 1 (423 K) ( ), and k cm 3 1 (473 K) ( ) molecules 1 s 1. An attempt was also made to measure k 1 above 473 K. A drop in k 1 of about % was observed at 523 K as compared to 473 K. As a check of the system, mixtures of 10 mbar of, 100 mbar of DME, and 5 bar of Ar was left for half an hour in the cell at 523 K and 296 K, respectively. The oxygen content of hot mixtures determined by mass spectroscopy was less than 1/5 of that in the cold mixture after half an hour. Consumption of oxygen by DME is therefore important at temperatures above 473 K making it impossible to measure k 1 at higher temperatures with the present experimental setup. In Figure 3(B), k 1 is plotted as a function of 1000/T. A straight line through the data in Figure 3(B) gives k 1 ( ) exp((46 27)/T) cm 3 molecule 1 s 1. Hoyermann et al. [17] obtained an expression for the temperature dependence of k 1 ; k 1 ( ) exp(( )/T) cm 3 molecule 1 s 1. Using this expression we obtain a value of k 1 which is 60% higher than the values reported by Masaki et al. [15]

7 OXIDATION OF DIMETHYL ETHER 633 dependence while the temperature dependence of k prod,0 is uncertain. It could be argued that k prod,0 may increase with temperature due to the additional energy deposited into the O # 2 complex at higher temperatures. Conversely, it could be argued that k prod,0 may decrease with increasing temperature, for entropy reasons, i.e., the tight transition state involved in the transfer of a hydrogen from the group to the CH 2 OO group in the OO # radical. Additional studies are important to confirm the temperature dependence of k prod,0 as reported by Hoyermann et al. [17]. Rate Constants for the Reaction First, k 5 was studied at 296 K and 1 bar total pressure using the low pressure system. Following radiolysis of gas mixtures of 5 mbar of DME and 995 mbar of SF 6 the transient absorption was recorded at 300 nm. The decay of the absorption transients were fitted using a second-order expression: Abs(t) (A 0 A inf )/(1 k 2nd (A 0 A inf )t) A inf (III) Figure 3 (A) Plot of the pseudo-first-order decay rate constants obtained from fits to experimental absorption transients obtained with 310 nm as detection wavelength using 18 bar total pressure and 296 K (filled circles), 353 K (filled triangles), 423 K (filled squares), and 473 K (filled triangles) plotted as function of the oxygen partial pressure. (B) Rate constants for the reaction of with plotted as function of 1000/T. The straight line is a linear regression through the data giving k 1 ( ) exp(( 46 27)/T ) cm 3 molecule 1 s 1. at room temperature. k 1 obtained by Hoyermann et al. [17] cannot be compared directly with the result obtained here. The value of k 1 reported by Hoyermann et al. [17] was obtained using 1 5 mbar total pressure and independent of total pressure in the pressure range applied. Hence, Hoyermann et al. [17] measured the low pressure value of k 1, k prod,0. It is interesting to note the negative temperature dependence of k prod,0 determined by Hoyerman et al. [17]. It is expected that k RO2,0 has a strong negative temperature where Abs(t) is the time dependent absorbance at 300 nm, and the three parameters; A 0, A inf, and k 2nd are the absorbance of radicals extrapolated back to time zero, the absorbance extrapolated to infinite time, and the second-order rate constant for reaction (5) times 2 ln10/( l), respectively. The absorption transients were always well fitted using expression (III). The reciprocal of the half-life derived from the fits to the data are plotted as a function of the maximum transient absorbance in Figure 4(A). As seen from the figure, 1/t 1/2 is proportional to the maximum transient absorbance. A linear regression through the data give a slope of ( ) 10 5 s 1. This slope equals 2k 5 ln10/( l). Using CH3OCH2(300 nm) cm 2 [6] and cm we obtain k cm 3 molecule 1 s 1 5 ( ) The uncertainty includes both uncertainty in the slope of the straight line in Figure 4(A) and 12% uncertainty in the absorption cross section of the radical. The linear regression of the data in Figure 4(A) gives a small positive intercept of ( ) 10 4 s 1. This intercept is probably due to a small amount of (only 0.07 mbar is necessary to explain the intercept) in SF 6 leading to a removal of radicals. To complete this study, k 5 was studied at 18 bar

8 634 SEHESTED ET AL. tions for the radical at 280, 300, and 310 nm of CH3OCH2 2.20, 2.74, and cm 2 [6], respectively. The obtained values of k 1 are plotted in Figure 4(B) as a function of the temperature. As seen from the figure there are no discernible effects on k 5 of the temperature in the studied temperature range. We therefore choose to determine an average of k 5 at all temperatures. The value of k 5 obtained from the high pressure experiments is ( ) cm 3 molecule 1 s 1. The uncertainty included both uncertainty in the data in Figure 4(A) and 12% uncertainty in the absorption cross section of the radical. Hoyermann et al. [17] report a value of k 5 of ( ) cm 3 molecule 1 s 1. This value is significantly lower than k 5 determined here. This could be explained as a pressure dependence on k 5 going from 18 bar to a few mbar total pressure. However, considering the number of degrees of freedom in ( ) # 2 which is presumably formed from reaction (5) it seems unlikely that there is a significant pressure effect on k 5. The reason for this apparent discrepancy is uncertain at present. Figure 4 (A) The reciprocal half life of the decay of the radical plotted as function of the maximum transient absorbance at 300 nm. The straight line through the data is a linear regression. (B) Rate constant for reaction (5) determined at 18 bar total pressure using the decay of the radical with 280, 300, or 310 nm as detection wavelengths. total pressure in the temperature range K. The rate constants were obtained by second-order fits with the experimental transients using 280, 300, or 310 nm as detection wavelength. To avoid any influence from the reaction of radicals with oxygen pres-ent in the SF 6, high purity SF 6 was used (see experimental for specifications) and the reaction mixture was irradiated with at least five electron pulses to remove any oxygen via reaction with radicals before the experiment. The rate constants, k 2nd kln10/( l), obtained fitting the absorption decays were converted to k 5 using the optical path length of 20 cm and the absorption cross sec- Rate Constants for the Reaction M 9: CH 2 O M When the temperature was increased above 573 K a sharp increase in the loss rate of the radical was observed. We ascribe this increased loss rate to thermal decomposition of radicals. In Figure 5(A) transient absorptions obtained at K following pulse radiolysis of mixtures of 100 mbar of DME and 18 bar of SF 6 are shown. The transients were fitted between s using expression (I). The transient absorption at the first 5 s was not used in the fit because of the overshoot of the photomultiplier due to the electron pulse. The pseudo-first-order rate constants obtained from the fits are plotted in the Figure 5(B). It is evident from Figure 5(B) that the decay rate increases with increasing temperature. The smooth line in Figure 5(B) is obtained from k exp( 12800/T ) s 1 as suggested by Loucks and Laidler [18,19] with an offset of 3.7 due to reaction (5) is plotted in Figure 5(B) for comparison. As seen from the figure the decay rate of radicals determined here is consistent with k 2 determined by Loucks and Laidler [18,19]. The decay rates of the radical obtained in this work is determined over a narrow temperature range, hence, we can not determine an independent rate constant for k 2. However, our work clearly supports the work of Loucks and Laidler [18,19] and confirms their expression for k 2.

9 OXIDATION OF DIMETHYL ETHER 635 M 9: products (5) k We derive the following rate constants for reaction (1) at 296 K: molecule 1 s 1, k prod,0 ( ) k RO2,0 ( ) cm 3 cm 6 molecule 2 s 1, and k RO2, ( ) cm 3 molecule 1 s 1. The temperature dependence of k RO2, was studied in the temperature range K and k RO2, ( ) exp((46 27)/T) cm 3 molecule 1 s 1 was determined. Values of k 5 ( ) cm 3 molecule 1 s 1 at 1 bar total pressure and a temperature of 296 K and k 5 ( ) cm 3 molecule 1 s 1 at 18 bar total pressure in the temperature range K was obtained. Finally, absolute measurements of k 2 in the temperature range K support exp( 12800/T ) s 1 2 as determined by Loucks and Laidler [18,19]. This and previous work have some implications concerning the combustion chemistry of DME. In this work we showed that k RO2, is relatively independent of temperature. It is expected that in the high pressure system the only product of reaction (1) is the peroxy radical,. Any OH radicals produced by reaction (1) would abstract a hydrogen atom rapidly via reaction (6): OH O 9: H 2 O (6) Figure 5 (A) Transient absorptions (310 nm) at 473 ( 0.025), 523 ( 0.02), 573 ( 0.015), 623 ( 0.010), 643 ( 0.005), and 666 K and 18 bar total pressure following irradiations of 100 mbar dimethyl ether and 18 bar SF 6. To separate the absorption transients the absorbance given in the parenthesis was added. The smooth lines are first-order fits to the experimental data. (B) Pseudo-first-order decay rate observed using a detection wavelength of 310 nm as a function of the temperature. k 2 determined by Loucks and Laidler [18,19] plus 3.7 (due to reaction (5)) are shown for comparison, see text for details. DISCUSSION In previous sections we present kinetic data on reactions (1), (2), and (5): This would lead to an apparent decrease of the loss rate of radicals. k 1 determined at 18 bar total pressure was relatively independent of temperature and, hence, the data showed no evidence for OH radical formation from reaction (1). At ignition temperatures the formation of peroxy radicals becomes reversible: M ;: M (1c, 1c) Hence, the importance of this channel of reaction (1) decreases. No evidence for reversibility of reaction (1) was encountered in this work. At higher temperatures the second channel of reaction (1) may be important: 9: 2HCHO OH (1b) M 9: products M 9: CH 2 O M (1) (2) This reaction may be an important ignition process of DME. For example to model the chemical reactions

10 636 SEHESTED ET AL. in the combustion of DME, Dagaut et al. [9] used k exp(337/t ) cm 3 molecule 1 s 1 prod In this context it is interesting to note the temperature dependence of k prod determined by Hoyermann et al. [17], k prod ( ) exp(( )/T ) cm 3 molecule 1 s 1. The value for k prod used by Dagaut et al. [9] is more than 20 times lower at all temperatures than the value determined by Hoyermann et al. [17]. However, the temperature dependence of the rate constant, k prod, determined at low pressures by Hoyermann et al. [17] is not necessarily representative for k prod (T) at higher pressures where # the energy distribution in the excited radicals is changed rapidly by collisions. However, if k prod determined by Hoyermann et al. [17] is representative at high pressures and high temperatures the value of k 1 used by Dagaut et al. [9] must be increased by a factor of more than 20. This clearly increases the importance of reaction (1) in the model proposed by Dagaut et al. [9] for the oxidation of DME at K suggesting that reaction (1) may play a role in the ignition of DME. Clearly, a study of k prod at high temperatures (above the temperature range used in this study) and high pressures is desirable. Also an absolute study to confirm the value of k 2 determined by Loucks and Laidler [18,19] is needed. It should be mentioned that reaction (2) is expected to be the most important loss reaction for radicals at combustion temperatures above 1000 K [9]. We would like to thank the Danish Technical Research Council and Ford Motor Company for financial support of this project and T. J. Wallington and E. W. Kaiser, Ford Motor Company, for helpful discussions. BIBLIOGRAPHY 1. M. A. Roubi, Chem. & Eng., 44, 37 (1995). 2. T. Fleisch, C. McCarthy, A. Basu, C. Udovich, P. Charbonneau, W. Slodowski, S.-E. Mikkelsen, and J. Mc- Candless, Soc. Automotive Eng., , 1 (1995). 3. S. M. Japar, T. J. Wallington, J. F. O. Richert, and J. C. Ball, Int. J. Chem. Kinet. 22, 1257 (1990). 4. M. E. Jenkin, G. D. Hayman, T. J. Wallington, M. D. Hurley, J. C. Ball, O. J. Nielsen, and T. Ellermann, J. Phys. Chem., 97, (1993). 5. T. J. Wallington, M. D. Hurley, J. C. Ball, and M. E. Jenkin, Chem. Phys. Lett., 211, 41 (1993). 6. S. Langer, E. Ljungstrøm, T. Ellermann, O. J. Nielsen, and J. Sehested, Chem. Phys. Lett., 240, 53 (1995). 7. J. Sehested, T. E. Møgelberg, T. J. Wallington, E. W. Kaiser, and O. J. Nielsen, J. Phys. Chem., (1996). 8. K. J. Hughes, P. D. Lightfoot, and M. J. Pilling, Chem. Phys. Lett., 191, 581 (1992). 9. P. Dagaut, J. C. Boettner, and M. Cathonnet, Twenty Sixth Int. Sym. on Comb., Napoli, 28th July 2nd August, H. F. Calcote, Comb. Flame, 42, 215 (1981). 11. K. B. Hansen, R. Wilbrandt, and P. Pagsberg, Rev. Sci. Instr., 50, 1532 (1979). 12. J. Sehested, K. Sehested, O. J. Nielsen, and T. J. Wallington, J. Phys. Chem., 98, 6731 (1994). 13. J. Sehested, Risø-R-804, T. J. Wallington, P. Dagaut, and M. J. Kurylo, Chem. Rev., 92, 667 (1992). 15. A. Masaki, S. Tsunashima, and N. Washida, J. Phys. Chem., 99, (1995). 16. L. Nelson, O. V. Rattigan, R. Neavyn, and H. W. Sidebottom, Int. J. Chem. Kinet., 22, 1111 (1990). 17. K. Hoyermann and F. Nacke, Twenty Sixth Int. Sym. on Comb., Napoli, 28th July 2nd August, L. F. Loucks and K. J. Laidler, Can. J. Chem., 45, 2767 (1967). 19. L. F. Loucks and K. J. Laidler, Can. J. Chem., 45, 2763 (1967).