898
Setser et al.
arguments suggest that the polyatomic group acquires rotational energy and angular momentum. Acknowledgment. One of us (D.J.B.) wishes to acknowledge helpful discussions with Drs. D. S. Y. HSU,J. W. Hudgens, and W. A. Sanders. We are grateful for the support provided by the National Science Foundation
(MPS75-02793). References and Notes (1) D. J. Bogan and D. W. Setser, J . Chem. Phys., 64, 586 (1976). (2) K. C. Kim, D. W. Setser, andC. M. Bogan, J. Chem. Phys., 60, 1837 (1974). (3) (a) K. C. Kim and 0. W. Setser, J. Phys. Chem., 77, 2493 (1973); (b) J. H. Parker, Int. J. Chem. Klnet., 7 , 433 (1975). (4) R. L. Johnson, K. C. Kim, and D. W. Setser, J. Phys. Chem., 77, 2499 (1973). (5) (a) R. K. Soily and S. W. Benson, J. Am. Chem. Soc., 93, 1592 (1971); (b) D. M. Goiden and S. W. Benson, Chem. Rev., 69, 125 (1969). (6) W. H. Duewer and D. W. Setser, J. Chem. Phys., 58, 2310 (1973). (7) F. R. Cruickshank and S. W. Benson, Int. J . Chem. Klnet., 1, 381 (1969); Deo(H-CH20CH3)is 4-5 kcai mol-‘ less than DedH-CH2CH3) which implies the presence of a radical stabilization energy. (8) S. W. Benson, J. Chem. fduc., 42, 502 (1965). The OH single bond energy varies greatly from one compound to another but good explanations have not been advanced to explain these bond energy changes. (9) (a) R. B. Bernstein and R. D. Levine, Adv. At. Mol. Phys., 11, 216 (1975); (b) R. D. Levine and R. B. Bernstein, Acc. Chem. Res., 7, 393 (1974); (c) R. D. Levine and R. B. Bernstein, “Modern Theoretical Chemlstry”, W. H. Miiier, Ed., Plenum Press, New York, N.Y., 1975; (d) R. D. Levine, B. R. Johnson, and R. B. Bernstein, Chem. Phys. Lett., 19, 1 (1973). (10) M. A. Nazar, J. C. Polanyi, W. J. Shrlac, and I. J. Sioan, Chem. Phys., 16, 411 (1976). (11) D. E. Mann, B. A. Thrush, R. D. Lide, J. J. Ball, and N. Acquista, J. Chem. Phys., 34, 420 (1966). (12) (a) J. M. Herbelin and G. Emanual, J . Chem. Phys., 60, 689 (1974); (b) R. Herman, R. W. Rothery, and R. J. Rubiin, J. Mol. Spectrosc., 2 , 369 (1958). (13) H. E. O’Neai and S. W. Benson, “Thermochemistry of Free Radicals”, in “Free Radicals”, J. K. Kochi, Ed., Wiiey-Interscience, New York,
N.Y., 1973, pp 275-359. (14) W. A. Chupka and J. Berkowitz, J . Chem. fhys., 54,5126 (1971). (15) B. de B. Darwent, Natl. Stand. Ref. Data Ser., Natl. Bur. Stand., No. 31 (1970). (16) K. Tamagake and D. W. Setser, unpublished results. Experiments with an improved arrested relaxation reaction vessel and with observation of the emission intensity with a Fourier transform spectrometer have confirmed the F CH30CH3distributions of Flgure 3. (17) (a) J. C. Pohnyi and K. 6. Woodail, J. Chem. Phys., 56, 1563 (1972); (b) A. M. G. Ding and J. C. Polanyi, Chem. Phys., 10, 39 (1975). (18) (a) D. H. Mayiotte, J. C. Polanyi, and K. 6. Woodall, J. Chem. Phys., 57, 1547 (1972); (b) C. A. Parr, J. C. Poianyi, and W. H. Wong, lbkl., 58, 5 (1973). (19) (a) I. Procaccia, Y. Shlmoni, and R. D. Levine, J. Chem. Phys., 83, 3181 (1975); (b) R. D. Levine, R. 6. Bemstein, P. Kahma, I. Procaccla, and E. T. Upchurch, ibld., 64, 796 (1976). (20) Double peaked rotational distributions normally are associated with two different reaction dynamics, see the following references for examples: (a) H. Heydtmannand J. C. Polanyi, App. Opt., 10, 1738 (1971); (b) M. A. Nazar, J. C. Polanyi, and W. T. Skrlac, Chem. Php. Lett., 29, 473 (1974). (21) (a) D. W. Smith and L. Andrews, J . Chem. Phys., 60, 81 (1974); (b) G. E. Ewing, W. E. Thompson, and G. C. Pimentei, /bid., 32, 927 (1960); (c) D. E. Miiiigan and M. E. Jacox, ibid., 51, 227 (1969). (22) T. Shimanouchi, Natl. Stand. Ref. Data Ser., Natl. Bur. Stand.,No. 39 (1972). (23) (a) A. M. G.Ding, L. J. Kirsch, D. S. Perry, J. C. Polanyi, and J. K. Schreiber, Faraday Discuss., Chem. Soc., 55, 252 (1973); (b) K. Tamagake and D. W. Setser, unpublished resuks. These data agree well with the results from ref 23a. (24) J. A. Austin, D. H. Levy, C. A. Gottiieb, and H. E. Radford, J. Chem. Phys., 60, 207 (1974). (25) G. Herzberg, “Molecular Spectra and Molecular Structure, 111. E W o n Spectra and Electron Structure of Polyatomic Molecules”, Van Nostrand, New York, N.Y., 1966, p 612. (26) H. W. Chang, D. W. Setser, and M. J. Perona, J. Phys. Chem., 75, 2070 (1971). (27) Y. Beers and C. J. Howard, J . Chem. Phys., 84, 1541 (1976). (28) R. L. Redlngton, W. B. Olson, and P. C. Cross, J. Chem. Phys., 36, 1311 (1962). (29) R. Foon and M. Kaufman, Prog. React. Klnet., 8, No. 2, 81 (1975). (30) D. J. Smith, D. W. SeW, K. C. Klm, and D. J. Bcgan, J. Phys. Chem., following paper in this issue. (31) J. P. Sung and D. W. Setser, Chem. Phys. Left., in press.
+
HF Infrared Chemiluminescence. Relative Rate Constants for Hydrogen Abstraction from Hydrocarbons, Substituted Methanes, and Inorganic Hydrides D. J. Smith,“ D. W. Setser,” K. C. Kim,lb and
D. J.
Bogan’’
Department of Chemistry, Kansas State University, Manhattan, Kansas 66506 (Received November 1 I, 1976)
The relative HF infrared emission intensities from the reactions of F atoms with 24 organic and inorganic hydride molecules, RH, have been measured in a flowing-afterglowapparatus at room temperature. By operating under conditions such that the emission intensity is first order in [RH] and the relaxation of the initial HF’ vibrational populations was minimal, rate constants, relative to CH4,for HF formation were obtained. By combining these measurements with absolute rate constants for selected members of the series, the relative rate constants can be converted to absolute room temperature rate constants. Since the initial HFt vibrational populations have been measured for these hydride molecules, the current measurements can provide absolute rate constants for formation of individual HFt vibrational levels. In addition to the relative rate constant measurements from the flowing-afterglow apparatus, chemiluminescence results from cold-wall arrested relaxation experiments are presented for reactions of F + various ethers and methanol-dl. The HFtUJdistributions from $he cyclic ethers and the magnitude of the rate constants for ethers vs. alkanes are discussed with respect to the enhanced rotational energy disposal for the F + CH30CH3reaction.
The Journal of Physlcal Chemistry, Vol. 8 1 , No. 9 , 1977
HF Infrared Chemiluminescence
899
been observed in a room temperature flowing afterglow apparatus to obtain relative rate constants for H F formation. The technique will be discussed and results will be presented for some hydrocarbons, substituted methanes, CH30CH3, CH3COCH3,and some inorganic hydrides. These relative rate constants can be combined with absolute measurements for some member15 of the series and with initial HFt population datal-'' to obtain absolute rate constants for formation of HFt in individual vibrational levels. The reactions which yield obviously extended HFt rotational populations, e.g., SiH4, GeH4, PH3, and CH30CH3,are especially interesting because of the possibility that these rotational populations reflect larger reaction cross sections relative to cases which do not yield extended populations. Relative to the HF distributions from most other hydrocarbons, the rotational distribution from F + (CH3)20are especially puzzling; therefore, the cold-wall, arrested relaxation studies were extended to include F + cyclic ethers, DOCH3,and Si(OCH3I4and the initial HFt vibrational-rotational distributions from these reactions also are reported here. The dagger superscript will be used to denote initial HF, populations; if this aspect is not important the dagger usually will be omitted.
Experimental Section The relative rate constant measurements were made with a conventional flowing afterglow apparatus, using Ar as the carrier gas, pumped with a large mechanical pump and a Roots-type blower. The carrier gas was passed through a molecular sieve trap cooled to -77 "C and was metered with a precalibrated Fischer-Porter floating ball flow meter. The linear flow velocity in the 5 cm diameter cylindrical aluminum flow reactor was 85 m/s. The carrier gas flow was divided into two parts. About 15% was mixed with a flow of CF4 (0.3% of the total carrier gas flow) and passed through a microwave discharge (80 W) cavity to produce F atoms; the remainder passed directly to the flow tube. The discharge section of the apparatus was an alumina tube, which lasted the duration of the study without developing leaks. The alumina tube extended 8 cm into the flow tube and terminated in a 2-cm diameter Pyrex bulb (attached via epoxy) containing numerous small holes. The hydride reagent, premixed with argon, was added to the flow 8 cm further downstream from the terminus of the alumina tube (F atom inlet). Observations of H F emission were made through NaCl windows, fitted to the aluminum flow tube by a O-ring seal, The center of the first viewing window was 4 cm downstream from the reagent mixing port, which was a 3-cm diameter glass ring constructed from 3-mm Pyrex tubing containing numerous small holes punched on the inside of the ring. The only important calibration needed for relative emission measurements was the flow meter used to measure flows of the RH reagent. Since the reagent mixtures were typically 1/100 to 1/500 of RH in Ar, the flow meter was calibrated for argon by measurement of pressure drop from standard volumes. All the flow measurements for the RH mixtures were made with the same flow meter. The mixtures were prepared by measuring the pressure of reagent (less than 10 Torr) with a Barocel electronic manometer (a transducer type pressure gage) and then adding Ar to the reservoir to obtain the desired pressure; the argon pressure was measured with a mercury manometer. The HF emission intensities were observed with the same monochromator and detection system that was used for the arrested relaxation studies, see below. Spectra initially were taken of the Au = 1 bands with reduced slit width in order to obtain sufficient resolution to assign rotational distributions. At 1Torr these distribution
-
% (f l
$2
\ I
3200
-
\ I
3400
3600
WAVENUMBER
38C0
4200
4000
CM-'
+
Figure 1. Low resolution HF emission spectra from F CH4, CPHB, and SiH,. The simulated spectra, shown as dotted lines, are displaced upward for convenience of display. These spectra were obtained from the high [F] conditions of table; the vibrational distributions are given in Tables I and 11.
were found to be room temperature Boltzmann even for reagents such as SiH4 and GeH4 which yield extended initial rotational distribution^.^ After it was established that the rotational distributions were Boltzmann, low resolution scans were sufficient to establish relative vibrational populations via computer simulation of the low-resolution spectra, see Figure 1. These low-resolution scans were taken for each reagent and they were used to provide the relationship between the signal level at a given wavelength and the total area of the spectrum. For the rate constant measurements, plots of intensity (at a single wavelength setting, which was selected to be the peak corresponding to the overlap of the 1-0 and 2-1 bands) vs. RH flow were obtained and compared to that of CHI for the same concentration of F atoms, see Figure 2. The ratio of the slopes of these linear plots of intensity vs. concentration of RH or CH4,adjusted for the total relative areas associated with the Au = 1emission, gave the relative rate constants for formation of HFt(u I 1). Further consideration of the method is given in the Results section. The reaction vessel and the associated experimental techniques used for the arrested relaxation experiments have been described previously."" The reaction reported here (methanol, ethylene oxide, dioxane, trioxane, and tetramethoxysilane) were studied concurrently with the previously reported results from the F + cyclic alkanes5 and were done prior to the F + RHCO series described in the preceeding paper.' However, the only difference in technique is that the atoms were produced by microwave discharge of SFG in a quartz, rather than an alumina, tube. Before the introduction of reagents (standard commercial grade), the whole apparatus, including the glass handling Torr. section, was routinely pumped to better than Except for trioxane and ethylene oxide, the vapor pressure of the liquid at room temperature waa used as the reservoir The Journal of Physlcal Chemistry, Vol. 81, No. 9 , 1977
Setser et al.
900
Scheme I 100
k
FtRH&HF(u=O)t
CGH12
R
?Al-,?kgllRHl PkF, tF1 k
F t RH+HF(u=
1) t R
' ? A , - , ? ~ Q ~ I R H I' ? k ~ ~ [ F l k
FtRH+HF(u=2)tR ?A,-,?kg,[RHl
?k~~[Fl
k
F t RH-",HF(u=3)tR
Flgure 2. A typical plot of HF emission intensity vs. concentratlon of hydride reagent. The intensity is the signal at a fied wavelength, whlch was the maximum from the overlap of the P-branch of the 1-0 band and the R-branch of the 2-1 band. The different RH plots have been normalized so as to correspond to the CH4 plot shown here. A flow of 1.0 X mol s-' corresponds to a concentration of 4 X 10'' molecules ~ m - ~ .
pressure. For trioxane, the reagents flask was maintained at 60 "C to provide the necessary flow rate into the reactor; ethylene oxide was metered from a reservoir at 100 Torr pressure. Concentric mixing nozzles made from quartz were used for these studies. A crossed nozzle geometry in which the F atom flow was crossed with flow of substrate also was tried for a few experiments with tetramethyoxysilane; the results were the same as with the concentric nozzle. The typical flows were 4 pmol/s of SF6 and 10 pmol/s for the reagent; the measured pressure in the Torr, cold-wall reactor was -1 X All spectra were recorded with a Perkin-Elmer Model 210-B monochromator fitted with a liquid N2cooled PbS detector. The signal was amplified with a Princeton Applied Research phase-lock amplifier. The fundamental spectra, 4300-3100 cm-l, were taken with a 0.4-mm slit and the overtone spectra, 8000-5000 cm-l, were obtained with a 2.0-mm slit. Relative HFtd populations were deduced from computer simulation2i6of the observed spectra. Both the fundamental and overtone spectra were used; however, greater emphasis was placed on the fundamental spectra. Since representative examples of the HF emission spectra were shown in the preceeding paper, only populations distributions will be presented here. The previous arguments2p5 claiming that vibrational relaxation is fully arrested for most reagents, but that rotational relaxation is only partly arrested, apply to the experiments with methanol-dl, ethylene oxide, dioxane, trioxane, and tetramethoxysilane.
Results A. Relative Rate Constants for R H with Primary C-H Bonds. The reaction kinetics will be examined to show how relative rate constants were obtained from the emission data. The reactions with primary C-H bonds giving HFt distributions similar to those from CH4will be considered in this section. There is little direct formation of HFt(u = 0 ) for these compounds3and, in particular, we will assume k(u = 0 ) is negligible for CH4. The most important HFt relaxation processes are collisions with RH and F. Vibrational relaxation by the argon carrier gas is of no importance, providing impurities are controlled, and the concentration of HF,is too low for bimolecular processes involving HF, to be i m ~ 0 r t a n t . l ~We also have omitted diffusion to and quenching at the walls in Scheme I. The values of (u = 1-4) are 188.6, 319.8, 398.3, The Journal of Physical Chemistry, Voi. 81, No. 9 , 1977
and 429.7 s-l, respectively.20 Kinetic modeling of this scheme by computer integration of the rate equations with vibrational relaxation rate constants21i22 from the literature (RH = CH4) showed that under the experimental conditions a nonsteady state kinetics prevailed, Le., the total concentration of HF, increased with distance along the tube. Depending upon the concentration of F and RH and the reaction time, vibrational relaxation from the initial HFt, distribution can be important. In order to simply relate the total observed intensity -
-
-
-
which includes emission from u = 1up to the maximum level (urn), to formation rate constants, the intensity, I", should reflect the formation rate into each vibrational level rather than relaxation processes. The emission intensity also should be f i t order in [RH]. Considerable effort w a ~ expended to discover such experimental conditions. The most important factors are low [F] and [RH], efficient mixing, 1Torr pressure of carrier gas, and observation as close as practical to the point of mixing. In this work the observation window was -4 cm (0.4 ms) from the mixing ring. The concentrations used for most experiments with the substituted methanes were [F] E 0.5 X 1013 and [RH] = 1-400 X lo1' molecule cmS. This estimate for [F] is based upon an assumedz350% conversion of CF4 to 2F + CF2. This [F] concentration was adopted after a series of exploratory experiments. With further experience it was learned that [F] could be lowered by a factor of 10 and satisfactory emission intensity still could be obtained. Therefore, several experiments also were done with reduced [F], see footnote b of Tables I and 11. Since vibrational relaxation of HF, is faster by most RH than by F, it is not necessarily advantageous to have [RH] > [F]. However, for the second set of conditions [RH] was not increased as CF4was reduced; thus, the net effect was to reduce relaxation since F and CF, fragments were reduced and [RH] was constant. Comparison of the initial HFt, distributions (from arrested relaxation experiments)2i3 to the distributions of Table I shows: (i) the distributions for the reduced [F] conditions are virtually identical with those from arrested relaxation experiments and (ii) the main effect of the higher [F] is a higher u1/u2 ratio for some, but not all, cases. Within experimental error, the same relative rate constants were obtained for both sets of conditions. The familar integrated form of the rate equation for a simple bimolecular reaction is
-
where [RH], and [F], are the initial concentrations of RH and F, and [HF] is the concentration of HF at time t.
HF Infrared Chemiluminescence
9Ot
TABLE I: Relative" Rate Constants for Organic RH RH
Lit. values
This work"
Vibrational distribution n.:n,:n,
CH,
1.0
1.0
H,
0.39 i O . O l e 0.34 i O.05f 0.42.f 0.20g 0.41' 2.9 i 0.2e 3.4.f 0.3f 2.0' 5.4 i 0.2e 6.5 f 0.6'
0.37 0.36b
0.35 :0.53:0.12 (0.22:0.66:0.12)b 0.45:0.45 :0.10 (0.29:0.56:0.15)b
2.1 3.0b
0.35:0.45:0.20 (0.18:0.50:0.32)b
CZH, C(CH3 1 4 (CH3
4.5 2.7 1.6 1.6b 0.60 0.15 (0.19)c O.OSb ( 0.09)c 1.6 (2.4)c
)2O
(CH3)2C0
CH,CHO CH,Cl CH,CN CH,NO, c-C,H, 6'-'
0.40 i 0.5g
-
3.2 i 0.007e 1.5' 4.3 f. 0.4e
2
CZH, 6'
H6
CH, F
0.58 f 0.02e 0.94h 0.11 f 0.0V 0.20h 0.010 f 0.002e 0.003h 0.004 i 0.002g
CHJ, CHF,
0.40:0.50:0.10 0.35:0.43:0.22 0.33:0.50:0.17 (0.26:0.33:0.34:0.07)b 0.35:0.43 :O. 22 0.45 :0.40 :0.1 5 (0.77:0.20:0.03)b 0.60:0.40: 5 for all u levels. Direct com arison of data from the same experimental conditions2 show that the dimethyl ether reaction gives larger populations in higher J levels than does the CH30H reaction. If the truncation procedure suggested by ref 11 is followed, estimates of initial rotational distributions indicated by the dotted lines of Figure 3 are obtained. These estimates for the cyclic ethers give ( f ~=) 0.10; corresponding estimates for the cyclic alkanes were ( f R ) = 0.07. The main difference is that the maxima in the estimated initial distributions from Figure 3 occur at 1-2 higher J values than for the cyclic alakanee. The rotational distributions for HzCO or (CH3)ZO extend to higher levels and the maxima occur at 3-4 higher J values than for the distributions of Figure 3. However, the available energy is larger than for the cyclic ethers and ( f ~for ) H2C0 or (CH3)2O is only increased to -0.12. The principal conclusion is that the mere presence of an oxygen atom adjacent to the C-H abstraction site does not ensure a high rotational energy disposal.
r
Discusfiion A. Abstraction Rate Constants. The results in Tables I and I1 demonstrate that measurement of infrared chemiluminescence in a fast flow system is capable of giving good values for relative rate constants, providing that precautions are taken to identify the optimum operating conditions. Comparison of the measured and literature values of the relative rate constants for the more thoroughly studied cases, such as Hz, C2H6, and C(CHJ4, shows that the agreement is excellent. For most other cases our results tend to agree better with the smaller ratios from the literature, although the trend is barely noticeable. With the exception of CHBF2and CHF3 (vide infra), the uncertainty in kRH/kCH4 is estimated to be &15% based upon the reproducibility of the rate constant ratios obtained from different experiments. With RH compounds for which relaxation is especially severe or for which a high direct yield of HFf(u = 0) occurs, the method is less satisfactory. Extension of the fast-flow chemiluminescence method to measurement of relative rate constants of other reactions should be feasible. As demonstrated by the work of ref 22, it also is possible to study vibrational relaxation of €ITt,in F + RH flowing-afterglowreaction systems. The data of Tables I and 11,as well as that of ref 10, show that the technique can give initial vibrational distributions, providing that proper precautions are taken. For this purpose the method is somewhat less reliable than the arrested relaxation method; however, it is more versatile and can be applied to more complicated chemical systems. One general problem was encountered which merits discussion. With nearly all reagents, the plot of IT vs. RH showed downward curvature at sufficiently high flows of RH. It was shown with several reagents that the effect was not related to HFt, vibrational relaxation. Other possibilities may be flow dependent mixing patterns or deThe Journal of Physical Chemistry, Vol. 819 No. 9 , 1977
pletion of [F] so that eq 6 does not hold. In some cases, competing reactions, such as addition to double bonds or abstraction of h a l ~ g e n , l ~ "can " ~ ~occur and removal of F without formation of HF can be severe. Even for arrested relaxation experiment^,^ we have experienced considerable difficulty in obtaining reproducible HF+,J populations from reactions of F with highly substituted halomethanes; the relative vibrational populations appear to be sensitive to flow rates and possibly nozzle geometry. In the present work, the HFt, populations from CHzF2and CHF3 are shifted toward u = 1, relative to the distribution expected on the basis of the available energy. Since high concentrations of CHzFzand CHF3, - 5 X 10l2,were required in order to observe the emission, there may have been some HFtvrelaxation. Also, the intensity vs. [RH] plots were especially curved for CHzF2and CHF3 Although our work shows that H atom abstraction rate from CH2Fzand CHF3 is rather slow, no quantitative information about the abstraction rate constant should be inferred. Other techniques also have demonstrated that the F atom removal rate by CHzFzand CHF3 is slow, relative to CH4 The most recent study31 cm3 with CHF3 reported a rate constant of 1.5 X molecule-ls-' for F atom removal at room temperature, which is in moderately good agreement with the work of cm3 moleClyne, McKenney, and Walker14 (2.2 X cule-' s-'). The apparent agreement between the rate constants for F atom removal and HF formation (Table I) from F + CHF3 should be regarded with caution, because of our curved intensity vs. CHF3concentration plot. In our opinion further work is necessary to identify reaction products before the F + CHF3 reaction is adopted as a standard for H atom abstraction measurements." The possibility of secondary or unusual competing reactions32 should not be overlooked. The reaction with CH3N02also is unusually slow and required a concentration of 5 X 10l2 molecule cm-3 to obtain a spectrum. In this case, arrested relaxation data" show that the u = 1 population is larger than for u = 2. Thus, vibrational relaxation may not have seriously affected the data of Table I. Nevertheless, direct formation and/or relaxation to HF(u = 0) may have occurred and the relative rate constant for CH3N02in Table I probably is a lower limit. In contrast to CH2Fzor CHF3, the plot of intensity vs. [CH3N02]was linear with [CH3N02]. The dipole moments33of CH3Cland CH3Fare 1.8 D and the dipole moments of CH3CN and CH3N02are 3.9 and 3.5 D. The decline of the H abstraction rate constant with increasing dipole moment is quite noticeable. The reduction in abstraction rate constant for CH3CN is even more surprising since the bond energy34is 5 kcal mol-l less than for C2H6. Evidently a substituent on the methyl group that reduces the electron density on the hydrogen atoms (hence giving a large dipole moment) reduces the abstraction rate constant. It is surprising that the abstraction rate constants for C2H435and C6H6are nearly as large as for CHI in spite of the competition provided by the addition channels, which are -3 times faster than the a b ~ t r a c t i o n .Addition ~~ to the unsaturated portion of the aldehyde and ketone molecules also probably occurs, but these abstraction reactions also are comparable to those for analogous alkanes. Based upon the relative yield of HF(v = 4) from CH3CH0, which arises from abstraction at the aldehyde position, and the magnitude of the total relative rate constant, the rate of abstraction of the aldehyde hydrogen seems to differ little from the abstraction rate with primary C-H bonds. N
HF Infrared Chemiluminescence
The reaction rate of F with SiH4 and GeH4 are nearly equal, but -6 times faster than with CHI. The recommended15rate constant for CH4at room temperature is 8.0 X cm3 molecule-’ s-l; therefore, the rate constants for SiH4 and GeH4 are -5 X lo-’’ cm3 molecule-’ s-’. The reason that the rate constant for GeH4 is not larger than for SiH4may just be that both are nearly equal to the gas kinetic value. The rates of reaction with HBr and HI also are comparable, but -7 times faster than for HCI. There is generally good agreement in the literature regarding the slow rate constant for F + HCl. Within the limit of the combined experimental errors, our data agree with those of Jonathan and co-workers for the hydrogen halides, although our results give a ratio of 1.1rather than 1.4 for HI vs. HBr. These changes in the HC1, HBr, HI, CH4, SiH4, GeH4 series presumably relate to changes in the activation energies, which is between 1 and 1.5 kcal mol-’ for CH4.I5The rates for H atom reactions with CHI, SiH4,and GeH4continue to increase throughout the series and kGeH4/kSiH4 B. Rotational Energy Disposal and Magnitudes of Rate Constants. The larger reactive cross sections of SiH4and GeH4 and the greater available energy, relative to CHBX compounds, may explain the enhanced rotational energy disposal for the F + SiH4 and GeH4 reactions. The rate constants for CzH6and (CH3)20are nearly the same, but the rotational distributions differ.2~~ Thus, the magnitudes of the reactive cross section cannot be invoked to explain the difference in the HFt rotational distributions from these reactions. Evidently the presence of an oxygen atom adjacent to a CH3 group does not ensure formation of HFt, in high rotational states either since the rotational distributions from CH30H were rather normal. Nevertheless, the data from Si(OCH3)4and trioxane do suggest that either rotational relaxation is more easily arrested in these systems or that a higher fraction of the available energy is released to ER(HF)for these particular ethers. Further workz9to establish the initial HFtvJrotational populations for a series of reagents is required before these apparent differences in rotational energy disposal can be explained. Acknowledgment. The aid of Mr. J. P. Sung in the completion of Tables I and I1 is greatly appreciated. This work was supported by the National Science Foundation Research Grant (MPS 75-02793) and by a NATO Research Grant (No. 739).
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References and Notes (1) (a) Current address: Queen Mary College, London, United Kingdom. (b) Current address: Los Alamos Scientific Laboratory, Los Alamos, N.M. (c) Current address: Code 6180, Naval Research Laboratory, Washington, D.C. (2) D. J. Bogan, D. W. Setser, and J. P. Sung, J. Phys. Chem.,preceding paper in this issue. (3) D. J. Bogan and D. W. Setser, J. Chem. Phys., 64, 586 (1976).
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The Journal of Physlcal Chemlstty, Voi. 81, No. 9, 1977
