J. Pkys. Ckem. 1995, 99, 16284-16289
16284
Rate Coefficients of the Reactions of C2H with 02, C2H2, and H20 between 295 and 450 K Hilde Van Look and Jozef Peeters" Department of Ckemistn, Universiq of Leuuen, Celestijnenlaan 200F, B-3001 Leuuen, Belgium Received: May 22, 1995; In Final Form: July 19, 1995@
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The absolute rate coefficients of three of the most important C,H-removal reactions in hydrocarbon flames, products (rl), C2H CrHr products (r2), C2H H20 products (r3), have been measured C2H 0 2 over the temperature range 295 < T < 450 K by a laser photodissociationkhemiluminescence (LPD/CL) technique. Ethynyl radicals were generated by 193 nm excimer laser photolysis of C2H2, and their real-time pseudo-first-order decay was monitored by the CH(ArAd-X2n) chemiluminescence resulting from their reaction with 0 2 , present in a high and constant concentration. It was ascertained that in the experimental conditions of the decay experiments, at a pressure of I O Torr He, the C2H radicals were electronically and vibrationally relaxed. The results for the first two reactions can be expressed by the Arrhenius relations kl = (1.9 & 0.1) exp[(O&IO)/T(K)] cm3 molecule-' s-l, in close x lo-" exp[(+160&15)/T(K)] and k. = (1.3 0.2) x agreement with the more recent literature data. For reaction r3, which was investigated for the first time, our data can be fitted to the expression kj = (1.9 i 0.2) x lo-" exp[(-200&30)/T(K)] cm3 molecule-' s-l. The reaction between C2H H2O appears to be quite fast and should therefore be incorporated in hydrocarbon combustion models.
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Introduction
Gas phase reactions of ethynyl radicals are of fundamental chemical interest and occur in a wide variety of natural and artificial environments. These radicals have been detected in interstellar space and in planetary atmospheres.' In the combustion of fossil fuels, ethynyl radicals are involved in the formation of polyacetylenes and possibly of polycyclic aromatic hydrocarbons (PAH);'.? they may also intervene in NO, chemistry. The elucidation of the role of C2H radicals in combustion processes requires the knowledge of the rate constants of the principal formation and destruction reactions. The mechanism of C2H formation in hydrocarbon combustion is not yet fully understood. All recent determinations and ab initio calculations of the C-H bond energy in C?H? agree on a high value of ~ 1 3 1 .kcal/mol;4,5 5 it follows that H-abstraction from ClH2 by H. 0, and OH will face energy barriers of at least 28.5. 30, and 13 kcal/mol, respectively, such that even the most favorable process can only be of marginal importance. In this laboratory, evidence was obtained that C:H radicals arise in C?H?/O/H atomic flame systems by a long reaction sequence initiated by 0 reaction, involving CH, and CH as the primary C2H2 intermediates, and ending with 0-attack on C3H2.h The present work focuses on the reactions of C2H with 0 2 , ClH?, and HzO, which are all important CzH-removal reactions in hydrocarbon flames. Several groups have investigated the kinetics of the reactions of C2H with 0 2 and C2Hr (rl and r2). The earlier literature data show a scatter of a factor 3, even at 300 K, mainly as a result of the fact that ethynyl radicals are difficult to monitor. Recently, more reliable methods to follow CrH radicals were developed, and the resulting kinetic data show a better accord. All previous measurements of the rate constants of the C2H 0, ( r l ) and C,H C2H2 (r2) reactions at room temperature are summarized in Table 1. Lange and Wagner' used a dischargelflow apparatus with mass-spectrometric detection. Their rate constants were only lower limits because of incomplete mixing in their observation
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Abrtract published
in
Adttrntr ACS Ah\truct\
October 1. 1995
0022-3654/95/2099- 16284$09.00/0
TABLE 1: Literature Values of the Rate Constants of the C2H O2 and C2H C2Hz Reactions at 295-300 K, in cm3 molecule-' s-l k(C?H + C?H?) k(C2H + 0:) .5.5 x IO-"" >5.O x IO-" " Ldnge and Wagner ( 1975)3.1 x lo-" Laufer and Bazs (1975y Renlund et al ( 198 I)" 2 x lo-" 5x Ldufer et a1 (1984)1'
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Stephens et a1 ( 1 9 8 7 ~ " Lander et a1 (1990)'? Shin and Michael (1991 j l ' Koshi et al (1992j2I Pedersen et a1 ( 1992) I' I(' Fdrhat et a1 ( 1991) I this work
1.5 x lo-"'
4.2 x lo-" 2.9 x 10-11
1.3 x lo-"' 1.5 x 1.3 x 1.9 x 1.3 x
3.3 x lo-" 3.3 x 10-1'
lo-"' lo-"' lo-'(' IO-"'
Value at 320 K
zone. Laufer and co-workers8.l0determined the rate constant for the C2H C2H2 reaction (r2) from the appearance rate of C4H2 in the flash photolysis of C2H2. The results obtained by Renlund et al.y,24were most probably the rate constants for the reactions of vibrationally and/or electronically excited ethynyl radicals. Stephens et a1.I' and Lander et a].'? measured the rate constants by following the time-decay of a C2H infrared transient absorption line using color center laser spectroscopy. The same method was used by Pedersen et al.I5.lhand by Farhat et al." Shin and Michael'j used atomic resonance absorption spectroscopy to monitor H atoms formed in the CrH ClH2 reaction. Koshi et a].?' followed the rate of the reaction with acetylene by measuring the appearance rate of CJH?. The main objective of the present work was to measure the rate constant of the reaction between C2H and H20, which has not yet been investigated to our knowledge. This reaction might proceed by exoergic H-abstraction:
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C,H
+ H,O - C2H2+ OH
Mor,,, = - 13 kcaVmol (r3a?
and could be fast enough to be a major C1H-loss mechanism in hydrocarbon flames. given their high H20 content. The reaction might (re)generate C2H2.
0 1995 American Chemical Society
Kinetics of C2H
+ 02, C2H2, H20 at 295-450
J. Phys. Chem., Vol. 99, No. 44, 1995 16285
K
pressure
1.20
flow controllers
1.oo
0.80 0.60
interference filter photomultiplier
thermocouple entmnce window c w i mer1ieht
0.40 0.20 0.00 4.0
6.0
8.0
10.0
12.0
14.0
16.0
reaction time (p)
optics laser
0.50 0.00
Figure 1. Schematic diagram of experimental apparatus. The excimer laser beam (not shown) is perpendicular to the luminescence collection optics axis.
Direct measurements of the rate coefficients of the title reactions, at temperatures between 295 and 45 1 K, were canied out using a laser photodissociatiotdchemiluminescencetechnique (LPDKL). In this technique, introduced by Wittig et al.,9,23.24 the time-resolved decay of [C2H] is monitored by the CH(A-X) chemiluminescence resulting from the reaction C2H 0 2 CH(A2A) C02 at large excess of [Oz]. In our study, we have paid particular attention to possible vibrational excitation of the ethynyl radicals, and we ascertained that the kinetic measurements were carried out in conditions of total pressure and gas composition ensuring complete relaxation.
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Experimental Section Apparatus. A schematic diagram of the experimental apparatus is given in Figure 1. It is part of a LPDLIF setup described C2H radicals were generated pulsewise in a blackened stainless-steel cell by photodissociation at 193 nm of C2H2 using the output of an ArF excimer laser (EMG 101 MSC, Lambda Physik) with the beam section telescoped down to 0.8 cm x 0.3 cm. The pulse length was -20 ns; the repetition frequency was 10 Hz; and the pulse energy was -30 mJ. The 430 nm CH(A2A-X211) chemiluminescence from the reaction region was detected with suitable collection optics. The spectral region of interest was isolated by a narrow bandpass interference filter (429.5 f 4.0 nm). The photomultiplieroutput was fed to an SRS type 250 boxcar averager. The delay between the laser pulse and the opening of the boxcar gate was gradually increased at a rate of 0.1 psis, while the gate width was kept constant at 1 ,us. Ethynyl decays were monitored over reaction times of 5-20 ,us. Data acquisition was achieved by a Macintosh II computer. The reaction-cell is resistively heatable up to a temperature of 450 K. The temperature is measured by a calibrated chromeV alumel thermocouple. The total pressure, measured by a Datametrics Barocel sensor, was in the 1-10 TOKrange; helium was used as the main bath gas. All gases were obtained commercially; the purities were He 99.995% (UCAR), C2H2 99.6% (UCAR), and 0 2 99.995% (L'Air Liquide). Since commercially available acetylene contains acetone, acetylene was prepurified before usage by passing it through a trap cooled by a dry-ice-acetone mixture. The flow rates of the gases were regulated and measured by calibrated mass flow controllers (MKS). The concentrations of the species were determined from the fractional flows and from the total pressure. The total gas flow was sufficiently large to refresh the gas in the observed volume between successive laser shots. Photolytic Generation of C f i . The 193 nm absorption cross section of the precursor C2H2 is about 2 x cm2;19hence,
-0.50 -1.00 -1.50
-2.00 -2.50
-3.00 -3.50
I
I'
4.0
""" "
I.".""'I
'. " . ' " ' I " . . ' 8.0
6.0
10.0
""
I ' "'""'I
12.0
""
14.0
.-.--I 16.0
reaction time (p) Figure 2. Typical decay and semilog plot of the relative C2H concentration vs reaction time; T = 366 K; ptoc= 10 Torr (He bath gas); [02] = 4.9 x lOI5 molecule/cm3;[CZH~] = 4.9 x 1014molecule/ cm3. The straight line represents the weighted linear least-squares tit.
at an excimer laser fluence of 1.2 x 1017 photons/cm2 pulse only -2.5% of the precursor is dissociated. The photochemistry of acetylene is a complex matter; H atoms, CH, C2, C2H, and triplet vinylidene (3B2) radicals were detected by several To avoid production of C2 and CH, which at 193 nm can occur only by multiphoton processes, the laser fluence was kept here below 120 mJ/cm2. The only photofragments from a 193 nm one-photon process are C2H and H. In acetylene photolysis at 193 nm, the ethynyl radicals are formed in high-lying vibrational levels of the X2C state in The latter lies only 3692 cm-' addition to the A211 (10.5 kcaYmol) above the X?X state. Shokoohi et a1.I8 have measured quenching constants for the A211 state by He: kq = cm3 molecule-' s-l. At a total pressure of 5 TOK 4 x He, the lifetime of C2H(A211) is only 1.3 ,us such that after 5 ,us the C2H radicals are 198% in the X21: state. The quenching of A211can result in vibrationally excited ethynyl X2C radicals. Method. The real-time exponential decay of ethynyl, at large excess of reactant(s), was monitored by the CH* chemiluminescence I(CH*) due to its reaction with 0 2 , present in large and constant concentration.
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C2H O2
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products
CH(A~A) + Co2
I
CH(X211)+ hv (430 nm)
The method is based on the quasi-steady-state for CH(A2A), which will be rapidly established and maintained on condition that the CH(A2A) lifetime is appreciably shorter than the C2H the chemical lifetime of lifetime. Since z(CH*) I0.5 C2H has to be larger than a few microseconds. The concentration of the reactant(s) is in large excess over [C2H] such that the kinetics are pseudo-first-order and the C2H decay is exponential. An example of a C2H exponential decay is shown in Figure 2. A C2H lifetime of a few microseconds, Le. appreciably longer than z(CH*), will ensure at the same time that any CH(A2A) that might have arisen in the photodissocia-" tion of C2H2 decays much f&ter than the steady-state CH* signal
Van Look and Peeters
16286 J. Phys. Chem., Vol. 99, No. 44, 1995
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from C?H 0 2 , such that the C2H-decay can be monitored over a time range (usually 5-20 ,us) where any initial [CH(A'A)] will already have vanished. It was observed that the excimer laser pulse causes a broadband transient signal (320-480 nm, decay time % 2 ps). independent of the presence of He carrier gas and/or 0 2 in the cell. The probable cause is fluorescence of one of the cell materials. To account for this, a "blank" transient, without 0 2 in the cell, was acquired each time and subtracted from the experimental decay trace. First-order decay constants (k') were measured at constant [O?]and variable reagent concentration, both in large excess. A plot of k' as a function of the reagent concentration [A] should yield a straight line with a slope equal to the rate constant of interest and an intercept equal to the 02 and precursor contribution:
4x
2x
0 1
02
0
06
03
10zl
10
08
12
14
( 1 0 ' ~m o ~ e c . / c m ~ )
Figure 3. Pseudo-first-order rate constants k' plotted vs [ 0 2 ] for the reaction of CIH - O? at T = 295 K:pLo,= I O Torr (He bath gas): [CzH!] = 4.9x 1014 moleculeicm'. The solid line is a weighted linear least-squares fit. of 5-20 ,us, depend on the He pressure, but only up to z 4 Torr. The k' values at lower pressures are always higher than the asymptotic values, which are reached once ptotexceeds 4 In the experimental conditions, the rate of removal out of the Torr. The absolute difference between k' at ptOt= 2 Torr and observed volume by diffusion and convection amounts to only the k' at plot? 5 Torr is largest for the lowest [Or] and becomes 500-2500 s-l 26 and is entirely negligible here. negligibly small for [O?]1 10l6molecule cm-?. These effects As mentioned above, the generated ClH(X'2) radicals may are ascribed to remaining vibrational excitation of the C?Hinitially arise in higher vibrational states. Remaining vibrational (X2X) radicals at low He pressure and low [O?];furthermore, excitation of the C2H radicals in a decay experiment can the quenching efficiency of 0 2 appears to be at least 1 order of influence the rate of their reactions with the reactant A, with magnitude higher than that of He. In any case, a He pressure 0 2 , and with C2H2. Hence, incomplete relaxation at the of about 5 Torr is sufficient to ensure a fully relaxed k', even at beginning of a decay (f = 5 ,us) would manifest itself by a the lowest [O,] of 1.6 x 10" molecule cm-'. An example of dependence of the rate of the decay (from t = 5 ,us up to 20 p s ) a k' vs [ 0 2 ] plot, at p,,,! = 10 Torr and T = 295 K, is shown in on the pressure of the He carrier gas, until that pressure is Figure 3. sufficiently high to ensure full relaxation (at t = 5 ps). Of The bimolecular rate constant kl(C2H 0 2 ) obtained from course, the degree of vibrational relaxation (at t = 5 pus) may the slopes of k' vs [ 0 2 ] plots at 295 K in the pressure range be strongly influenced also by the partial pressure of 02: it is 5-10 Torr He is (3.3 3~ 0.3) x lo-" cm3 molecule-' S - I . This well-known that potentially reactive species-such as 02 in this result is in excellent agreement with the value recently obtained case-can be highly efficient vibrational-energy acceptors. by Opansky et aI.li for C2H CX22(0,0,0))a t p = 10-100 Torr: kl = (3.3 k 0.3) x IO-" cm3 molecule-' s-l. Results and Discussion Our complete set of kl data for the temperature range from 295 to 450 K, obtained at I O Torr He, is displayed in Table 2. The reactants are always present in at least a 200-fold excess The statistical error is usually 2-5%; the total indicated errors, over C2H. Contributions from secondary or radical-radical of about lo%, include estimated systematic errors. The reactions can be neglected since the time between collisions of Arrhenius plot is shown in Figure 4. The rate coefficient shows C?H with the reactant is several orders of magnitude shorter a slight negative temperature dependence: ki(T) = (1.9 & 0.1) than the time between collisions of two ethynyl radicals. At x lo-" exp[(+160&15)/T] cm3molecule-' s-'. This Arrhenius 193 nm 0 2 absorbs weakly on the edge of the Schumannexpression agrees within 15% with the result obtained by Runge bands." With an absorption cross section of 0.3 x lo-'? Opansky et al.,15who derived kl(T) = (1.5 f 0.3) x lo-" cm-?, only -4 x of the 02 will be excited. exp[(230+36)/T] cm3 molecule-' s- over a temperature range C2H 0 2 (rl). Kinetic experiments were performed at an between 193 and 350 K and a t p = 10-100 Torr. Their results acetylene number density of (4.8 & 0.1) x 10l4molecule/cm? are also plotted in Figure 4. The slight negative temperature and at varying oxygen number densities, (1.6-14) x 10" dependence suggests that the reaction proceeds through a peroxy molecule/cm3, at total pressures between 1 and 10 Torr He and radical intermediate, HCCOO', which may either redissociate at 295 K. To obtain meaningful data. the pressure of the diluent back into reactants or isomerize and fragment into reaction gas must be sufficiently high so that the ethynyl radicals are products; the fraction of the initial adduct that redissociates vibrationally cold during a decay experiment. It was found that becomes more important at higher temperatures, and the net the first-order decay constants k', measured over the time range rate of product formation will decrease. An explanation along TABLE 2: Bimolecular Rate Constants k in cm3 molecule-' s-l, p = 10 Torr He C2H 0 2 C'H A C:H?
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295 332 366 393 448
(3.3 f 0.3) x (3.1 i 0.3) x (3.0 i0.3) x (2.9 0.3) x (2.7 0.3) x
**
lo-!' IO-" lo-'' IO-" lo-)!
295 332 366 393 448
**
(1.33 0.20) x (1.34 0.13) x (1.33 ~ 0 . 1 3 x) 11.31 =k0.13) x (1.34+0.131 x
lo-!!' IO-'" 10-'" 10-I"
295 295 332 35 I 385 131 1.5 I
(9.5 i 1.0) x 10-1' (9.7 i 1.0) x 10-1: ( 1 . 0 9 i 0 . 1 1 ) x lo-" ( 1 . 0 6 ~ 0 . 1 1xj I O - ' ' ( l . 0 9 & 0 . 1 1 )x 10-" (1.19 i 0.12) x 10-l) ( 1 . 2 6 i - 0 . 1 3 ) x lo-"
Kinetics of C2H
+ 02, C2H2, H20 at 295-450
1 k (cm3molec.~'s")
K
J. Phys. Chem., Vol. 99, No. 44, 1995 16287
10-10
I -11 10
I
2.0
2.5
3.0
d
3.5
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+ I
+ *
4.5
4.0
5.0
I 5.5
1000 KiT Figure 4. Arrhenius plot of the rate coefficient of the C2H 0 2 reaction: (-U-) this work; (+) Opansky et al.;ls (A) Lange and Wagner;? ( 0 )Renlund et al.;9 (0)Laufer at a1.;I0 ( x ) Stephens et al.;" (0)Lander et a1.I2
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these lines, based on ab initio calculations, was recently given by Carpenter30 for the analogous C2H3 0 2 reaction, which was also found to have a negative Arrhenius activation energy: k = (6.6 k 1.3) x exp[(125f50)/T(K)] cm3 molecule-' s - I . ~ ~ The ab initio study showed that the initial vinylperoxy radical first isomerizes to the dioxiranylmethyl radical, a threemembered ring comprising the a-carbon and the two 0 atoms. A similar pathway can be expected for the ethynyl reaction: n C2H 02 HC=COO HC-COO. One obvious set of products would be COz plus CH, with the latter arising for a small fraction in the excited 2A state. At pressures where C2H relaxation is complete, i.e. for ptot 2 5 Torr, our kl values are independent of pressure. Furthermore, in the temperature-overlap region, our data at p = 5 to 10 Torr coincide with those of Opansky et al.I5 obtained in the 10-100 Torr range. The probable reason why no pressure dependence is observed is the short unimolecular lifetime of the vibrationally excited HCICOOt intermediate. Given the low density of states associated with the well depth of only some 40 kcaYm01,~~ and considering the limited number of vibrational degrees of freedom, one can expect a unimolecular lifetime for redissociation of I s, such that collisional stabilization can become important only at pressures above atmospheric. The rate constant should then show an increase, to attain an asymptote once collisional stabilization dominates. C2H C2Hz (r2). Determinations of k2, at ptot= 10 Torr, were performed simultaneously with the measurements of k1(C2H 0 2 ) above; the slope of the linear K vs [OZ]plots yields kl(C2H +O2), while the intercept gives kz(C2H C2H2). Room temperature determinations of k2(C2H C2H2) were also made at ptot= 2 Torr (He bath gas) in the conventional way, from the slope of k' vs [C2H2] plots at a constant [02] of (6.5 or 8.1) x l O I 5 molecule/cm3, and with [C2H2] ranging from l O I 4 to lOI5 molecule/cm3. The results of the conventional procedure, at 2 Torr, gave k2 values identical to that obtained by the intercept method, at 10 Torr. The apparently facile vibrational relaxation here is attributed to the fairly high [OZ]. The k2(295 K) result, (1.3 f 0.2) x 1O-Io cm3 molecule-' s-', is in close agreement with the more recent literature values (see Table 1). Our full set of k2 data, in the 295-450 K range, listed in Table 2, shows no significant temperature dependence. The data of Shin and MichaelI3 over the wide 295-1475 K range exhibit a slight T-dependence: k2(T) = 3.02 x exp(-253/T) cm3 molecule-' s-'. Their result at room temperature is (1.3 f 0.3) x lo-'' cm3 molecule-' s-l, whereas for the T = 1226-1475 K range they report an average value of (2.5 f.0.7) x
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k (cm3molec."s")
0.0
1.0
2.0
3.0
4.0
5.0
6.0
1000 KiT Figure 5. Arrhenius plot of the rate coefficient of the C2H iC2H2 reaction: (-U-) this work; (0)Koshi et a1.;21(- -) Pedersen et a1.;16 (A)Farhat et al.;'? (- * -) Shin et a l . I 3 cm3 molecule-' s-'. However, as stated by the authors, the temperature variation is not at all certain since there is a substantial error. Koshi et alS2'determined the rate coefficient C2H2 reaction by two different experimental for the C2H techniques, LPD/MS at T = 298-438 K and LPD/ST (shock tube) at T = 1600-2200 K. They obtained k2 = (1.5 & 0.3) x cm3 molecule-' s-I, independent of temperature. Pedersen et al.I6 investigated the reaction over the temperature range 170-350 K. Their rate coefficient results, obtained at p = 10100 Torr, are nearly independent of temperature: k2(T) = (1.1 & 0.2) x exp[(28&20)/T] cm3 molecule-' s-I. Their 300 K result k2 = (1.3 f 0.2) x cm3 molecule-' s-I coincides with ours. The most recent literature results, by Farhat et al.,I7 can be represented by k2(T) = (1.0 & 0.1) x 1O-Io exp[(+180&2O)/fl cm3 molecule-' s-l for T = 295-854 K. The above mentioned data and the present results are shown in Figure 5. There appears to be no clear-cut evidence for a definite temperature dependence of k2 in the temperature range between 170 and 2200 K. The high rate and the absence of activation energy strongly suggest that the reaction follows an addition-dissociation mechanism proceeding through a shortlived vibrationally excited C4H3+intermediate that easily fragments into C4H2 H. C2H HzO (r3). This reaction has not been studied earlier. In our experiments, formation of OH radicals by photolysis of H20 can be excluded; the excited state responsible for dissociation is the unstable A 'BI state, and the A X transition has been observed only between 145 and 186 nm. To add water vapor, a fraction of the He stream was led over a distilled-water surface. The H20 flow was determined from the loss of weight of the liquid. The weight change measurement was started after the adsorption-desorption equilibrium was established and the [HzO] number density in the cell had become constant, as observed from successive 3 min measurements of the decay constant k'. Establishment of a constant [H20] took some 10 min. The repeatability of the weight loss measurements over 3 h periods was better than 3%. The H20 concentration was varied between 0 and 2.6 x lOI5 molecule/cm3;thus, the partial pressure of HzO was always far below the vapor pressure. Given the constancy of the gaseous H20 number density in the cell and the constancy of the H20 throughput, at temperatures well above the dew point both for the cell and the upstream feed line, there is no apparent reason for systematic errors of the gas phase [H20] as determined from the H20 throughput, the flows of the various gases (He, 0 2 , and C2H2), and the total gas pressure in the cell. The acetylene concentration was in the range (3.9-4.9) x l O I 4 molecule/cm3, the 0 2 concentration was between (3.1 and 5.0) x l O I 5 molecule/cm3, and the total
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Van Look and Peeters
16288 J. Pkys. Ckem., Vol. 99, No. 44, 1995
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clearly the fastest. In this respect, the C2H H.0 reaction appears to be an anomaly. With D(H-X) = 118 kcal/mol, it is expected to have a higher activation energy and to be much slower than C?H H?; instead the E, value is lower and the k(298 K) value is much higher. This does not at all support the above assumption that the reaction between ethynyl and water is a direct H-abstraction process. Moreover, the H20 molecule has particular electronic structure features that set it apart from the three other molecules considered here and that could enhance its reactivity toward C2H: the lone pairs on the 0 atom. One can suggest a role for (one of) them in the formation of an initial, loose 0- or Jr-complex that would facilely isomerize to one of the more stable structures H2C=C-OHL or HC=CHOH'. which lie 45-50 kcal/mol below the reactants.?' Their vibrational energy content would be largely sufficient for fast direct dissociation into H2CCO plus H, and C2H2 plus OH. respectively, o r i n the case of HC=CHOH'-for fast 1,2 or 1,3 H-migration and subsequent dissociation to H2CCO plus H.'" Overall, the route to ketene plus atomic hydrogen is the most exoergic (AH,. -35 kcallmol). It is worth noting that ketene and H are also produced in the lowpressure OH C?H? addition reaction.32 One must therefore admit that there is reason to doubt that the C2H H20 reaction (i) is a direct H-abstraction process and (ii) produces (only) C2H2 plus OH. It would therefore be rather premature to use our kinetic data to derive thermochemical or related information based on the C'H H20 C,H? OH equilibrium. One of the objectives of future work will be a product analysis of the CzH H20 reaction. In view of the high H20 content of hydrocarbon flames, the fairly fast reaction between C2H and HzO is undoubtedly an important C,H-loss mechanism in such combustion systems. Its importance relative to the major competing reactions with C2H2 and 02 will of course vary strongly with the position in a flame, because of the rapidly changing concentrations of these reactants. Typical mole fractions in the middle of the flame front of a strongly fuel-rich, near sooting, air-supported flame are x(O2) = 0.05. .x(C,H.) = 0.03, andx(H20) = 0.15. Thus, using the kl, k?, and k; expressions above, extrapolated to 1600 K, the rates of the C2H reactions with 0 2 , CrH2, and H20 are roughly in the ratio I :4:2.5. The reaction with H20 will become dominant in the postcombustion region, where the mole fraction of C:H? becomes 50.01. Prediction of the effects of the C2H H2O reaction on flame velocities and other (macroscopic) observables would require detailed flame modeling, which is of course beyond the scope of this work. If C;H2 and OH are major products, one of the obvious effects of the reaction would be to retard the net consumption of C?H: and hence. most probably, to increase the time available for the production of the aromatic precursors of PAHs and of soot.
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16x
IO
5
14x105
I
1
,
05
0
10
1 5
[HzO]
20
25
30
molec./cm3)
Figure 6. Pseudo-first-order rate constants k' plotted vs [H?O]for the reaction C?H HzO at T = 451 K: pto,= 10 Torr (He bath gas): [O?] = 3.9 x 10l5 moleculeicm' and [C>H>]= 3.9 x 1OIJ molecule/cm'. The solid line is a weighted linear least-squares fit.
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2x10"
I1
I
k (cm3molec."s")
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8~10.'~
7~10.'~
+
6~10.'~ 20
25
3 0
1000 KIT Figure 7. Arrhenius plot of the rate coefficient of the CIH reaction. at temperatures between 295 and 451 K.
35
H.0
pressure was 10 Torr (He bath gas). We have determined ki(CrH H20), in the conventional way, in a temperature region between 295 and 451 K. An example of a plot of k' as a function of [H20], for T = 451 K, is displayed in Figure 6; the slope of the line results in a ki(T = 451 K) value of (1.26 i 0.15) x IO-" cm3 molecule-' s-l; the experimental intercept is 1.5 x 10' s - ' . in excellent agreement with the expected value koJO?] f ~ C ~ H ~ [=C(1.6 ~ Hf~0.2) ] x loi S K I . Such internal consistency was observed at each temperature. The example of a k' vs [H20] plot in Figure 6 also confirms that the precision of the [HzO] determinations by the weight-loss method is better than 5 % . The set of k3 data is presented in Table 2. Figure 7 shows the Arrhenius plot; the temperature dependence can be expressed by k3(7) = (1.9 f 0.2) x lo-" exp[(-200f30)/7'j cm3 molecule-' s - I . Since this is the first measurement of the rate constant of the C2H H20 reaction, comparisons can be made only with analogous reactions. Assuming direct H-abstraction, similar processes are the reactions of C2H with H2, CHJ, and CzHs. For C2H H2, recently reported Arrhenius activation energies. Ea, for the 300-800 K range are between 2.0 and 2.5 kcaV mol, while the Arrhenius frequency factor, A, for the 300-500 K range is ~2 x lo-" cm3 molecule-' s-I: the room temperature rate constant is (4-5) x 10-'3..'7.2'Absolute rate CHd and C2H C2Hh have been coefficients of C2H measured only at room temperature; recently reported values are 3.0 x lo-'? and 3.6 x lo-" cm? molecule-' s-l, respectively.'? For H?, C b , and C2Hhthere is a Polanyi-Evans type correlation between the k(295 K) value and the H-X bond dissociation energy, D(H-X), which is equal to 103.3, 102.7. and 99 kcal/mol, respectively; the most exoergic reaction is
+
+
+
+
+
+
Conclusions
The LPD/CL technique, using the CH* chemiluminescence CO? CH(A?A) to follow CzH, was caused by C2H + 0 2 shown to be a valid and highly precise method, in the proper experimental conditions, for C?H(X'C) kinetics studies. The rate coefficient of the C2H 0 2 reaction was found to exhibit a slight negative temperature coefficient in the 295450 K range. in excellent agreement with the results of Opansky et al.Is in the 193-350 K region. The rate coefficient of C2H + C?H2 determined in the region 295-450 K was observed to be independent of temperature. The value obtained is in close agreement with other recent studies.
-
+
+
+ 0 2 , C2H2, H20 at 295-450 K The rate coefficient of the C2H + H20 reaction was measured
Kinetics of C2H
for the first time to our knowledge. The result in the 295-450 K range can be expressed by k3(T) = (1.9 f 0.2) x lo-” exp[(-200&30)/T] cm3 molecule-’ s-l. Comparison with rate constants and activation energies of other C2H reactions leads to doubt that C2H H20 is a direct H-abstraction reaction. In view of its high rate constant, the reaction is expected to be one of the principal CzH-loss processes in hydrocarbon flames.
+
Acknowledgment. The financial support of the Commission of the European Communities (Science Programme, Contract SCl*CT91-0637) and of the Belgian National Fund for Scientific Research (NFWO Contract 1.5437.92 N) is gratefully acknowledged. H.V.L. is indebted to the Flemisch Institute for Science and Technology (IWT) for granting her a doctoral fellowship. References and Notes (1) Tucker, K. D.; Kutner, M. L.; Thaddeus, P. Astrophys. J . 1974, 193, L 11. (2) Bonne, H.; Homann, K. H.; Wagner, H. Gg. Symp. (Int.) Combust. [Proc.] 1965, 10, 503. (3) Frenklach, M.; Clary, D. W.; Gardiner, W. C.; Stein, S. E. Symp. (Int.) Combust. [Proc.] 1984, 20, 887. (4) Ashfold, M. Private communication. (5) Andzelm, J.; Sosa, C.; Eades, R. A. J. Phys. Chem. 1993,97,46644669 and references therein. (6) (a) Boullart, W.; Devriendt, K.; Peeters, J. Submitted for publication. (b) Boullart, W. Ph. D. Thesis, University of Leuven, 1991. (7) Lange, W.; Wagner, H. Gg. Ber. Bunsen-Ges. Phys. Chem. 1975, 79 (2), 165-170. (8) Laufer, A. H.; Bass, A. M. J. Phys. Chem. 1979, 83, 310. (9) Renlund, A. M.; Shokoohi, F.; Reisler, H.; Wittig, C . Chem. Phys. Lett. 1981, 84 (2), 293-299. (10) Laufer, A. H.; Lechleider, R. J. Phys. Chem., 1984, 88, 66-68. (11) Stephens, J. W.; Hall, J. L.; Solka, H.; Yan, W. B.; Curl, R. F.; Glass, G. P. J. Phys. Chem. 1987, 91, 5740-5743.
J. Phys. Chem., Vol. 99, No. 44, 1995 16289 (12) Lander, D. R.; Unfried, K. G.; Glass, G. P.; Curl, R. F. J. Phys. Chem. 1990, 94, 7759-7763. (13) Shin, K. S.; Michael, J. V. J. Phys. Chem. 1991, 95, 5864-5869. (14) Koshi, M.; Nishida, N.; Matsui, H. J. Phys. Chem. 1992, 96,58755880. (15) Opansky, B. J.; Seakins, P. W.; Pedersen, J. 0. P.; Leone, S. R. J. Phys. Chem. 1993, 97, 8583-8589. (16) Pedersen, J. 0.;Opansky, B. J.; Leone, S. R. J. Phys. Chem. 1993, 97, 6822-6829. (17) Farhat, S. K.; Morter, C. L.; Glass, G. P. J. Phys. Chem. 1993, 97, 12789- 12792. (18) Shokoohi, F.; Watson, T. A.; Reisler, H.; Kong, F.; Renlund, A. M.; Wittig, C. J. Phys. Chem. 1986, 90, 5695-5700. (19) Satyapal, S.; Bersohn, R. J. Phys. Chem. 1991, 95, 8004-8006. (20) Okabe, H. J. Chem. Phys. 1975, 62, 2782. (21) Koshi, M.; Fukuda, K.; Kamiya, K.; Matsui, H. J. Phys. Chem. 1992, 96, 9839-9843. (22) Okabe, H. Photochemistry of Small Molecules; Wiley-Interscience; New York, 1978; p 177. (23) Reisler, H.; Mangir, M.; Wittig, C. Chem. Phys. 1980,47,49-58. (24) Renlund, A. M.; Shokoohi, F.; Reisler, H.; Wittig, C. J. Phys. Chem. 1982, 86, 4165-4170. (25) Peeters, J.; Vanhaelemeersch, S.;Van Hoeymissen, J.; Borms, R.; Vermeylen, D. J. Phys. Chem. 1989, 93, 3892-3894. (26) Peeters, J.; Van Hoeymissen, J.; Vanhaelemeersch, S.; Vermeylen, D. J. Phys. Chem. 1992, 96, 1257-1263. (27) Bauer, W.; Engelhardt, B.; Wiesen, P.; Becker, K. H. Chem. Phys. Lett. 1989, 158, 321-324. (28) Stull, D. R.; Prophet, H. Junuf Thermochemical Tables, 2nd ed.; NSRDS-NBS 37, 1971. (29) Melius, C. F. BAC-MP4 Heats of Formation and Free Energies; Sandia National Laboratories: Livermore, CA, 1988. (30) Carpenter, B. K. J . Phys. Chem. 1995, 99, 9801-9810. (3 1) Slagle, I. R.; Park, J.-Y.;Heaven, M. C.; Gutman, D. J. Am. Chem. SOC.1984, 106, 4356. (32) Kanofsky, J. R.; Lucas, D.; Pruss, F.; Gutmann, D. J. Phys. Chem. 1974, 78 (4), 311-316.
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