5050
J . Phys. Chem. 1987, 91, 5050-5054
-
ers,46-48a substantial increase of the medium viscosity obtained by cooling does not affect the quantum yields of the cis trans process while it strongly reduces the yields for the reverse reaction. The present findings confirm that when rotation about the -N=N- double bond is hindered, the photochemistry of azobenzene from the first and the second excited singlet state is the same. If in homogeneous media the photoisomerization occurs exclusively via the inversion mechanism (as suggested by Monti et a1.26)or via the rotation and inversion, according to the excited (46) Malkin, S.; Fischer, E. J . Phys. Chem. 1962, 66, 2482. (47) Gegiou, D.; Muszkat, K. A,; Fischer, E. J . Am. Chem. SOC.1968, 90, 12. (48) Gegiou, D.; Muszkat, K. A,; Fischer, E. J . Am. Chem. Sor. 1968,90, 3907.
state populated by light excitation (as suggested by R ~ U ~ ~is, ~ ’ ) , a question which, at the moment, remains unanswered.
Acknowledgment. This work has been financially supported by the ”Progetto Finalizzato per la Chimica Fine e Secondaria” of the C.N.R. The authors thank Prof. Bruno Samori of the Istituto di Chimica degli Intermedi, Facolta’ di Chimica Industriale, University of Bologna, for providing the JASCO-J 500 dichrograph and Mr. Gianfranco Gubellini for technical assistance. Registry No. a-CDx complex with trans-azobenzene, 109553-46-0; P-CDx complex with trans-azobenzene, 109553-47-I ; y-CDx complex with trans-azobenzene, 109553-48-2; a-CDx complex with cis-azobenzene, 109553-49-3;P-CDx complex with cis-azobenzene, 10955350-6; y-CDx complex with cis-azobenzene, 109553-5 1-7: azobenzene, 103-33-3;cyclodextrin, 126 19-70-4.
Flash Photolysis Resonance Fluorescence Investigatlon of the Gas-Phase Reactions of OH Radicals with a Series of Aliphatic Ketones over the Temperature Range 240-440 K Timothy J. Wallington and Michael J. Kurylo* Center for Chemical Physics, National Bureau of Standards, Gaithersburg, Maryland 20899 (Received: March 19, 1987; In Final Form: May 1 , 1987)
Absolute rate constants have been determined for the gas-phase reactions of OH radicals with a series of aliphatic ketones by the flash photolysis resonance fluorescence measurement technique. Experiments were performed over the temperature range 240-440 K at total pressures (using Ar diluent gas) between 25 and 50 Torr. The rate constant data for acetone ( k J , 2-butanone ( k 2 ) ,and 3-pentanone ( k , ) were used to derive the Arrhenius expressions k l = (1.7 f 0.4) X exp[-(600 f 7 5 ) / q cm3 molecule-‘s-I, k2 = (2.3 f 1.1) X exp[-(l70 f 120)/77 cm3 molecule-l s-l, and k3 = (2.8 f 0.3) X exp[(lO f 3 9 / 7 1 cm3 molecule-’ s-’. At 296 K, the measured rate constants (in units of cm3 molecule-I s-l) were k l = 2.16 f 0.16, k2 = 11.5 f 1.0, and k , = 27.4 f 1.3. Room temperature rate constants for the reaction of OH radicals with a number of other 2-ketones were also determined. These were (in the above units) 3,3-dimethyl-2-butanone (12.1 f O S ) , 2-pentanone (40.0 f 2.9), 2-hexanone (66.4 f 5.6), 2-heptanone (86.7 f 8.4), 2-octanone (1 10 f 9), 2-nonanone (1 22 f 13), and 2-decanone (1 32 f 12). The error limits represent 2 standard deviations (from the least-squares analysis); we estimate that an additional 5% uncertainty should be added to account for possible systematic error in the measurements. These results are discussed in terms of reactivity trends for C-H bonds located in the a , (3, and y positions with respect to the carbonyl group.
Introduction The gas-phase reaction with hydroxyl (OH) radicals represents an important loss process for ketones in the earth’s atmosphere.’S2 The need for a thorough understanding of the kinetics of these reactions is enhanced by the increased use of such oxygenates as fuel additives. Nevertheless, the extensive kinetic data base which exists for the reaction of OH radicals with many different classes of organic compounds3 does not extend to the reaction of OH radicals with ketones. More specifically, there are no studies on the temperature dependence of the rate constants; and multiple room temperature results exist only for acetone, 2-butanone, 4-methyl-2-pentanone, and 2,6-dimethyI-4-heptan0ne.~-” Fur(1) Finlayson-Pitts,B. J.; Pitts, Jr., J. N. Atmospheric Chemistry; Wiley: New York, 1986. (2) Carlier, P.: Hannachi, H.; Mouvier, G. Atmos. Enciron. 1986, 20,
2079.
(3) Atkinson, R. Chem. Rea. 1986, 86, 69. (4) Cox, R. A.; Derwent, R . G.; Williams, M. R. Enuiron. Sci. Technol. 1980, 14, 57. (5) Winer, A . M.: Lloyd, A. C.; Darnall, K. R.; Pitts, Jr., J. N. J . Phys. Chem. 1976, 80, 1635. (6) Cox, R. A.; Patrick, K. F.: Chant, S . A. Enoiron. Sci., Technol. 1981, 15, 587. (7) Zetzsch, C. 7th International Symposium on Gas Kinetics, Gottingen, Aug 1982. (8) Chiorboli, C.: Bignozzi, C. A,; Maldotti, A,; Giardini, P. F.: Rossi, A,; Carassiti. V . Int. J . Chem. Kine?. 1983, 15, 579.
thermore, for acetone and 2-butanone significant discrepancies exist in the reported room temperature data. We thus decided to conduct a systematic investigation of the kinetics of the reactions of OH radicals with series of ketones in order to resolve existing discrepancies in the literature and to provide a data base for the temperature dependence of these reactions. The kinetic measurement technique employed was flash photolysis resonance fluorescence (FPRF) at temperatures between 240 and 440 K and total pressures between 25 and 50 Torr (achieved with argon diluent). The results from these studies constitute the first measurements of the temperature dependence of the reactions of OH with acetone, 2-butanone, and 3-pentanone and the only existing rate constant measurements (at room temperature) for 2-heptanone, 2-octanone, 2-nonanone, 2-decanone, and 3,3-dimethyl-2-butanone.Our room temperature results for 2-pentanone and 2-hexanone are the first to be obtained by an absolute kinetic measurement technique. The series of 2-ketones was selected in order to develop a reactivity index (at 296 K) for C-H bonds positioned a,@, and y to the carbonyl group of the ketone. (9) Atkinson, R.; Aschmann, S. M.; Carter, W. P. L.; Pitts, Jr., J. N. Int. J . Chem. Kinet. 1982, 14, 839. (10) Edney, E. 0.;Kleindienst, T. E.; Corse, E. W . fnr. J . Chem. Kinet. 1986, 18, 1355. (11) Kerr, J. A.; Stocker, D. W. J . Atmos. Chem. 1986, 4, 253.
This article not subject to U.S. Copyright. Published 1987 by the American Chemical Society
OH Radical Reaction with Ketones
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987 5051
Experimental Section The apparatus and techniques used in the present work have been described p r e v i ~ u s l y . ~ ~Briefly, -'~ hydroxyl radicals were produced by the vacuum ultraviolet photolysis of H20(-0.1 Torr) by using a N2-pulsed discharge flash lamp equipped with Suprasil optics (A I165 nm). Following their production, the radicals were monitored as a function of time by the fluorescence (A2Z+ X2n,0-0 bond) excited by a microwave powered OH resonance lamp ( e l Torr of Ar saturated with water vapor). The signal was recorded with a cooled photomultiplier tube fitted with an interference filter (310-nm peak transmission, 10 nm fwhm) and positioned at right angles to both the flash and resonance lamps. Pulse-counting electronics and microprocessor-controlled multichannel scaling data storage comprised the data acquisition system. Depending on the initial O H concentration, fluorescence signal from 10 to 500 flashes was averaged to generate a kinetic decay curve suitable for least-squares analysis. Under pseudo-first-order kinetic conditions, reactant concentrations were chosen such that the first-order OH decay rates ranged from 50 to 1000 s-I and the O H radicals were monitored over at least three half-lives. Temperature regulation of the spherical, double-walled, Pyrex reaction vessel (=lo0 cm3 volume) was achieved by the passage of cooled methanol or heated oil between the outer walls. The gas temperature was measured by a chromel/alumel thermocouple projecting into the center of the reaction cell and was constant within 2 K during a set of experiments. Reaction mixtures consisting of water vapor, ketone, and Ar diluent were prepared manometrically in 5-L Pyrex bulbs and protected against photodissociation by room light prior to admission to the reaction cell. All pressure measurements were made with capacitance manometers and Bourdon gauges. In order to avoid the accumulation of photolysis or reaction products, experiments were conducted by flowing the gas mixtures through the cell with a typical residence time of 5 to 10 s. Since the flash lamp was operated at discharge energies of 40 to 125 J per flash with a repetition rate of 0.25 Hz, the reaction mixture was replenished every few flashes. In all cases the flash duration was negligible in comparison with the O H kinetic lifetimes. The argon diluent gas had a manufacturer's stated purity of 199.998% and was used without further purification. All of the ketones (with the exception of 2-decanone) had stated purities of at least 99%. The latter was supplied with a quoted purity of 98%. These reactants were subjected to repeated freeze-pumpthaw cycles prior to purification by a fractional distillation procedure in which the middle fraction was retained for use. For the apparatus setup and flash lamp conditions used in this investigation, the initial hydroxyl radical concentration was esbased upon a timated as 1Olo 5 [OH], I 10" molecules comparison with previous experiments from this laboratory and with similar systems used by other worker^.^^,^^ Thus, the reagent was high concentration range of (2-300) X l O I 3 molecules enough to assure pseudo-first-order kinetic conditions with respect to the radical decay. At these low radical concentrations, [OH] is directly proportional to the fluorescence signal and one can integrate the first-order rate expression to obtain
-
F, = F,, exp(-kIst(t - t o ) )= F,, exp(-(ko + k,[RH])(t - to)) (1) where F, and F,, are the O H radical fluorescence intensities at times t and to respectively, kist is the total first-order decay rate, ko is the first-order rate constant for O H removal in the absence of reactant (attributed to diffusion out of the viewing zone and reaction with possible impurities in the argon diluent gas), and k , is the bimolecular rate constant for the reaction of O H with (12) Kurylo, M. J.; Braun, W. Chem. Phys. Len. 1976, 37, 232. (13) Kurylo, M. J.; Cornett, K. D.; Murphy, J. L. J . Geophys. Res. 1982, 87, 3081. (14) Wallington, T. J.; Neuman, D. M.;Kurylo, M. J. Inf. J . Chem. Kinef., in press. (15) Wallington, T. J. Int. J . Chem. Kinet. 1986, 18, 487. (16) Witte, F.; Urbanik, E.; Zetzsch, C. J . Phys. Chem. 1986, 90, 3251
-
600
-
400
-
IVI
5 Y
I c
E v
200 -
0 40
20
0
60
80
100
P(ocetone), mtorr
Figure 1. Plot of (kist - k,) vs. acetone partial pressure at 296 K: (0) 25 Torr of Ar; ( 0 ) 37.5 Torr of Ar; (A) 50 Torr of Ar. The line represents a linear-least-squares analysis. 100,
.
I
1
I
.
!
,
I
1 '
2.0
,
,
1
2.5
3.0
3.5
4.0
4.5
1000/T, K-'
Figure 2. Arrhenius plots for the reactions of OH with acetone (0), 2-butanone ( 0 ) , and 3-pentanone (A).
reactant RH. Values of kist were determined by nonlinear least-squares exponential analysis of the experimental radical fluorescence decay curves. ko and k, could then be computed from the intercept and slope of a weighted linear least-squares fit of the kist values vs. reactant concentration. In practice, ko was measured experimentally (for the various conditions of pressure and flow rate) in the absence of reactant gas. Typical values of ko determined in this manner ranged from 20 to 40 s-' at T I 296 K (increasing to 40-80 s-l at elevated temperatures) and were constant to within 10 s-I for a given set of experimental conditions. Reactant concentrations were chosen such that ko was the minor component of kist and k, was thus determined from linear least-squares analysis of plots of (kist - k,) vs. [RH].
Results In all experiments, exponential decays of the resonance fluorescence signal were observed over at least three half-lives and the measured decay rates demonstrated a linear dependence on the ketone concentration. Figure 1, which shows the 296 K data for acetone, indicates the typical data statistics as well as the lack of a discernable effect of pressure on k , over the 25-50-Torr Ar pressure range. O H + CH3COCH3 products (1)
-
Similar observations were made for the rate data obtained with all of the ketones studied in this work. The linear least-squares analysis of the 296 K acetone data is indicated by the line drawn in Figure 1 and yields a value for the bimolecular rate constant of k l = (2.16 i 0.16) X cm3 molecule-' s-l where the error limits represent two standard deviations from the regression analysis. Rate constants for all of the OH + ketone reactions measured in this investigation are given in Table I. Variation of the flash energy (and hence [OH],) by a factor of 3 and of the reaction cell residence time by a factor of 4 resulted in no change in any of the rate constants measured, suggesting that the experiments
5052 The Journal of Physical Chemistry, Vol, 91, No. 19, 1987
W-allington and Kurylo
TABLE I: Rate Constants for the Reactions of OH with Ketones Measured during the Present Work k," cm3 molecule-' s-I reactant T = 240 K T = 296 K T = 350 K T = 400 K acetone 1.45 f 0.15 2.16 f 0.16 2.92 f 0.23 4.07 f 0.30 12.3 f 1.0 2-butanone 11.5 f 1.0 14.1 f 0.9 15.5 f 0.7 3,3-dimethyl-2-butanone 12.1 f 0.5
40.0 f 2.9
2-pentanone 3-pentanone 2-hexanone 2-heptanone 2-octanone 2-nonanone 2-decanone
28.5 f 1.7
27.4 f 1.3
66.4 f 5.6 86.7 f 8.4 110 f 9
29.1
* 1.7
27.9 f 3.2
T = 440 K
4.36 f 0.50 16.5 f 0.9 27.8 f 4.0
122 f 13 132 f 12
"The error limits are two standard deviations obtained by the least-squares analysis of the data. TABLE II: Comparison of the Present Results with Literature Data for the Reactions of OH Radicals with Ketones at 296 K k, cm3 reactant molecule-' s-I technisue" workers ref Cox et al. 4 55.3 RR acetone 7 2.3 f 0.3 FPRF Zetzsch Chiorboli et al. 8 RR 6.2 f 0.9
2-butanone
2.6 & 0.8 2.16 $z 0.16 35 f 10 9.5 f 0.9 12f2
9.7 f 1.7 11.5 f 1.0
3,3-dimethyl-2- 12.1 f 0.5 butanone 46.4 f 1.4 2-pentanone 40.0 f 2.9 18.2 f 3.3 3-pentanone 21.4 f 1.3 2-hexanone 89.7 f 6.0 66.4 f 5.6 86.7 f 8.4 2-heptanone 110 f 9 2-octanone 122 f 13 2-nonanone 132 f 12 2-decanone
RR FPRF RR RR FPRF RR FPRF FPRF
Kerr and Sheppard 11 this work 5 Winer et al. Cox et al. 6 Zetzsch I Edney et al. 10 this work this work
RR FPRF RR FPRF RR FPRF FPRF FPRF FPRF FPRF
Atkinson et al. this work Atkinson et al. this work Atkinson et al. this work this work this work
9 9
9
-
products
(2)
products
(3)
as a function of temperature from 240 to 440 K and are shown in Arrhenius plots in Figure 2. Linear-least-squares analyses of these data are indicated by the lines drawn in the figure and are represented by the following Arrhenius expressions (in units of cm3 molecule-I s-l):
k , = (1.7 f 0.4) k 2 = (2.3 f 1.1)
X
k3 = (2.8 f 0.3)
exp[-(600 f 75)/71 exp[-(170 f 120)/71
X X
R
"Calculated as 0.5kl (298 K). *Calculated as 0.5k3 (298 K). TABLE IV: Calculated CHI and CH2 Reactivities (with OH) in Aliphatic Ketones at 296 K reaction rate, io-" cm3 molecule-I sd
position to C=O
CH3
01
1.1 3.1
Y
this work
were free from complications due to secondary reactions involving either reactant photofragments or reaction products. The rate constants measured for reactions 1-3 were measured
+ CH3COC2HS O H + C2H5COC2H,
reaction rate, io-" cm3 molecule-I s-l in CH,(CO)R in CH3R 1.1Q 1.4 10.4 10.4, 13.8* 23.9 38.9 39.0 65.3 54.4 85.6 70.6 108.9 85.8 120.9 98.6 130.9
P
this work
"RR = relative rate, FPRF = flash photolysis resonance fluorescence.
OH
TABLE 111: Reactivity of Linear Aliphatic Chains with Respect to OH at 296 K
exp[(lO f 3 5 ) / g
The error limits are two standard deviations from the statistical analysis. Discussion The room temperature rate constants for all of the reactions of the present study are compared with literature data in Table 11. With the exception of the absolute measurements by Zetzsch' (also using FPRF), all of these comparisons are with relative rate investigations. Excellent agreement is observed between the present k , value and all of the existing literature values with the
53.7"
CHI 9.0 226
__-
"Assumed (see text). exception of the relative rate result reported by Chiorboli et a1.* Similarly, the rate constant for the 2-butanone reaction is in good agreement with the only absolute determination' and the two most recent relative rate studies.6si0 The origins of the discrepancy between all of these recent measurements and the high relative rate measurement of Winer et aLs is unclear. Comparison of our rate constants for the reaction of O H with 2-pentanone, 3-pentanone, and 2-hexanone with the relative determinations by Atkinson et aL9 shows agreement for 2-pentanone but approximate 30% differences for 3-pentanone and 2-hexanone. However, the error limits given in Table I1 represent only the random errors of measurement and do not include any estimation of possible systematic errors. While it is difficult to quantify systematic errors, we estimate a maximum of 5-10% for such contributions in the present work and somewhat larger uncertainties in the relative rate technique due to uncertainties in the reference rate constant. Thus, for these three reactions, agreement probably exists within the total combined experimental uncertainties. The Arrhenius preexponential factors and activation energies determined in this work for reactions 1-3 are somewhat lower than those for the corresponding OH alkane reaction^.^ Nevertheless (similar to the alkane reactions), a mechanism involving hydrogen atom abstraction from the ketones has been inferred from earlier product6 and kinetic' investigations. This matter will be considered in the ensuing discussion. Given the extensive data set on the reactivity of ketones with hydroxyl radicals presented in Table I, we are able to draw some interesting general conclusions about structure/reactivity relationships. To a first approximation, the rate constant for the
+
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987 5053
O H Radical Reaction with Ketones 2-butanone reaction (k,) appears to be (within experimental error) the arithmetic mean of the rate constants for acetone and 3pentanone over the complete temperature range of study. This is best seen in Figure 2. Such behavior suggests that the reactivities of the aliphatic chains (methyl and/or ethyl) on either side of the carbonyl group are essentially independent and additive. We can cm3 molecule-’ thus assign a rate of reaction (in units of s-l) for OH with a CH3group positioned a to the carbonyl group equal to one-half of k , or ~ 1 . at 1 296 K. Similarly, the reactivity of an ethyl chain adjacent to the ketone carbonyl can be taken as one-half of k3 or can be calculated by subtracting the C H 3 reactivity from k,. Reaction rates for OH with propyl, butyl, etc. chains adjacent to the ketone carbonyl can then be calculated by subtracting the ketone a-CH3reactivity from the rate constants measured for 2-pentanone, 2-hexanone, etc., respectively. These values are presented in Table 111. Using a similar analysis, subtraction of the ketone a-CH, reactivity from the rate constant for 3,3-dimethyl-2-butanone and dividing by 3, yields an approximate value for the rate of reaction of OH radicals with a methyl group positioned /3 to the carbonyl group of 23.7. Assuming that this value is also applicable to 6-CH, groups in straight-chain ketones such as 2-butanone and 3-pentanone, one can then calculate a value of =9 for the reactivity of an a-CH2 in the ketones. Thus, we find that the reactivities calculated for both the a-CH, and a-CH2groups in an aliphatic ketone are indistinguishable from the reactivities for CH3 and CH, groups calculated from O H kinetic data for normal alkanes.” Similar reactivities can also be expected for CH3 and CH, groups which are far removed from the carbonyl group in a ketone. Consequently, it seems reasonable to assume the P-CH, reactivity as an upper limit for the y-CH, reactivity in an aliphatic ketone. Under this assumption, the data for 2-pentanone can be used to compute a value of 226 for the reactivity of a P-CH, group. These group reactivity values are summarized in Table IV. It should be emphasized that the results of all of the above computations are highly approximate since they require taking the differences in large numbers which themselves carry experimental uncertainty. Also, the errors in such calculations are cumulative. The level of uncertainty can be demonstrated by the derivation of the a-CH, reactivity from our data on 2-butanone and 3-pentanone. Using the values first calculated for a- and P-CH3groups, one obtains reactivity values per P-CH, of 6.7 and 10.1 from these two ketones, respectively. Due to the presence of twice as many a-CH, groups in 3-pentanone as in 2-butanone, we decided to use a weighted average from the two compounds to determine the a-CH2 reactivity value of 2 9 reported above. Despite this degree of uncertainty, it is instructive to compare these reactivity assignments on a relative scale. Such a comparison reveals a significant enhancement of the group reactivity toward O H for both CH, and CH, when moved from the a to p positions with respect to the ketone carbonyl group. In both cases, the reactivity increases by approximately a factor of three. A similar rate enhancement has been observed by Atkinson et al.9 for aand P-CH, groups from a more limited sampling of aliphatic ketones. This increase is not expected from bond energy considerations since C-H bond energies are not expected to differ significantly from the a to the P positions.I8 However, it should be noted that a factor of three in reaction rate ate 298 K translates to a AE of only 650 cal which is well below the 1-2 kcal uncertainty levels of such bond energy tabulations. Also included in Table 111 are aliphatic chain reactivities calculated from rate constant data3 for O H with normal alkanes. For these computations, the CH, reactivity was set equal to one-half of the room temperature rate constant for OH C2H6. Values for the other alkyl chains (C2H5,C3H7,etc.) could then be calculated by subtracting the CH3value from the rate constants for C,Hs, C4H,o,etc. The increased aliphatic chain reactivities
+
(17) Atkinson, R.; Aschmann, S. M.; Carter, W. P. L.; Winer, A. M.; Pitts, Jr., J. N. Int. J . Chem. Kinet. 1982, 14, 7 8 1 . (18) McMillen, D. F.; Golden, D. M. Annu. Reu. Phys. Chem. 1982, 33, 493.
I
-
150 I
3
I
g
o
‘0)
CC
1
O/
2
4
6
Carbon Number
4
8
10
+
Figure 3. Straight-chain aliphatic reactivities obtained from OH ketone (0)and OH alkane (A)systems plotted vs. the carbon number of the chain. The lines are discussed in the text.
+
observed for C2 and higher in the ketones may be correlated with the enhanced (and possible y) CH, and CH, group reaction rates as shown in Figure 3. Here we have plotted the aliphatic chain reactivities calculated from both the ketone and alkane data vs. the carbon number of the chain. The line drawn through the alkane data is a linear-least-squares fit to this data set. The line through the ketone data is a smooth curve interpolation of the data. As can be seen, the methyl group reactivities from both data sets are indistinguishable. The slope of the line going through the C2 through C4 ketone data is significantly larger than that for the corresponding alkane data, whereas beyond C4 both data sets are essentially parallel. This latter observation suggests that the incremental increase in alkyl reactivity with the addition of a CH2 group to n-butyl and larger is, as expected, the same in both the ketones and alkanes. The larger incremental reactivities for the alkyl chains in the ketones are isolated to the first occurrence of P- and/or y-CH, and/or CH, groups (Le., for ethyl, propyl, and butyl). One possible explanation for these effects is the existence of a second reaction pathway in the OH + ketone systems in addition to the straightforward concerted hydrogen atom abstraction mechanism which is solely operative in the reaction of OH radicals with alkanes. This mechanistic alternative may involve coordination of the attacking OH to the carbonyl group to form a short-lived adduct which either decomposes back to reactants or forms products via H-form abstraction (preferentially with the 0, and possible y,substituents). In the case of P substituents, the formation of a six-membered ring could facilitate reaction.
Y
Along a similar vein, it is not unreasonable to suspect formation of a seven-membered ring, involving y substituents, thereby enhancing their reactivity as well. The occurrence of such an adduct formation channel is consistent with the observed trend of decreasing activation energies in reaction sequences l through 3. For both reactions 2 and 3, the activation energies are somewhat lower than those measured for the corresponding alkanes (propane and butane). Such coordination of the attacking OH radical with the carbonyl group is not unexpected in light of the electrophilic nature of the hydroxyl radical and the electron density within the carbonyl group. Similar effects have been suggested as influencing the reactive behavior of OH in other system^^^*^^ as well as the reactivity of other electrophilic species such as N032‘and O(3P).z2 (19) Wallington, T. J.; Atkinson, R.; Tuazon, E. C.; Aschmann, S. M. Inr. J . Chem. Kinet. 1986, 18, 837. (20) Kurylo, M. J.; Knable, G. L. J . Phys. Chem. 1984, 88, 3305.
J . Phys. Chem. 1987, 91, 5054-5057
5054
Finally, it is interesting to note that the room temperature gas-phase rate constants reported, herein for acetone, 2-butanone, 2-pentanone, and 3-pentanone are in good agreement with solution phasevalues (1.6 X lO-I3, 15 X lO-I3, 32 X lO-I3, and 23 X cm3molecule-I s-l respectively) measured by Adams et al.23 Such (21) Wallington, T. J.; Atkinson, R.; Winer, A. M.; Pitts, Jr., J. N. J . Phys. Chem. 1986. 90. 5393. (22) Cvetanovic, R. J.; Singleton, D. L.; Irwin, R. S. J . Am. Chem. SOC. 1981, 103, 3530.
agreement might be indicative of close similarities in the reaction mechanism for both phases.
Acknowledgment. The research described herein was conducted with the partial support of the National Aeronautics and Space Administration (Agreement W-15,816). (23) Adams, G. E.; Boag, J. W.; Currant, J.; Michael, B. D. 'Absolute Rate Constants for the Reaction of the Hydroxyl Radical with Organic Compounds" In Pulse Radiolysis, Ebert, M., Keene, J. P., Swallow, A. J., Baxendale, J. H., Eds.; Academic: New York, 1965; p 61.
A Mass Spectrometric Study of Mercury-Photosensitized Reactions of Monosilane and Ammonia Mixtures: Observation of Silylated Amines Ching-Hsong Wu Chemistry Department, Research Staff. Ford Motor Co.. Dearborn, Michigan 481 21 (Received: April 6, 1987)
Gas-phase stable and transient products in the Hg(3P1)-photosensitizedreactions of monosilane and ammonia mixtures in a static reactor were investigated by means of a mass spectrometer. Several Si-N-containing intermediates including silylamine (SiH3NH2),disilazane ((SiH,),NH), and disilaneamine (Si2HSNH2)were detected. The reactivities of these compounds were studied. Under UV irradiation these amines were found to be unstable and underwent reactions which presumably led to the formation of polymeric solids and silicon nitride. In the dark the half-life of silylamine at room temperature varied from 1 to 20 h depending on the surface activity of the reactor.
Introduction The feasibility of low-temperature deposition of silicon nitride (Si3N4) films on microelectronic devices using the mercury (Hg(3P,))-photosensitized reactions of monosilane (SiH4) and ammonia (NH,) mixtures has been recently dem~nstrated.'-~ Although the overall stoichiometry for the reactions can be simply represented by the equation 3SiH4
+ 4NH3 2 Si3N4(s) + 12H2 Hg
(1)
the chemical pathways by which the silicon nitride film is formed are rather complex. In this system, the chemical and physical processes are expected to consist of not only those occurring in the Hg-photosensitized reactions of the individual systems of SiH4 and N H 3 but also the intersystem reactions caused by radicalradical and radical-molecule chain processes in the gas and solid phases. Fore more than six decades, numerous studies have been conducted on the reaction mechansism of the Hg(3P,)-photosensitized reactions involving SiH44-Ioor NH3.l1-l9 However, ( I ) Peters, J. W. US. Patent 431 587, 1983. (2) Peters, J. W.; Gebhart, F. L.; Hall, T. C. Int. J. Hybrid Microelectron. 1980, 2, 59; Solid State Technol. 1980, 23, 121. (3) Padmanabhan, R.; Miller, B. J. J . Vac. Sci. Technol., A 1986, 4, 363. (4) Emeleus, H. J.; Stewart, K. Trans. Faraday SOC.1936, 34, 1577. (5) White, D. G.; Rochow, E. G. J . Am. Chem. Sac. 1954, 76, 3897. (6) Niki, H.; Mains, G . J. J . Phys. Chem. 1964, 68, 304. (7) Nay, M. A.; Woodall, G. N. C.; Strausz, 0.P.; Gunning, H. E. J . A m . Chem. Sac. 1965, 87, 179. (8) Varma, R.; Ray, A. K.; Sahey, B. K. Inorg. Nucl. Chem. Lett. 1969, 5, 497. (9) Kamaratos, E.; Lampe, F. W. J . Phys. Chem. 1970, 7 4 , 2267. (IO) Austin, E. H.; Lampe, F. W. J . Phys. Chem. 1976, 80, 2811. (11) Dickinson, R.; Mitchell, A. Proc. Natl. Acad. Sei. U.S.A. 1926, 12, 692. (12) Dickinson, R.; Mitchell, A. J . A m . Chem. SOC.1927, 49, 1478.
0022-3654/87/2091-5054$01.50/0
to date no mechanistic study of the reactions between silane and ammonia has been reported. In particular, very little is known about the formation and the subsequent reactions of the siliconnitrogen compounds. One of the simplest silicon-nitrogen compounds, which has been the subject of great theoretical interest in recent years, is silylamine (SiH3NH2). Many of its physical and chemical properties, such as the molecular structure,20-22the nuclear spin-spin coupling constant,23the ionization p ~ t e n t i a l ?and ~ the electron and proton a f f i n i t i e ~ ,have ~~,~ been ~ evaluated by quantum mechanical calculation. However, none of the these calculated values have been experimentally determined because the sample compound is not available. To our knowledge, silylamine has not been previously identified, and its existence is unknown. The existence of silylamine was indirectly inferred by Stock and Somieski.*' They reported that when chlorosilane reacted with an excess of ammonia there was some evidence for the formation of partially silylated amine products via the following reactions. (13) (14) (15) (16) (17) (18) (19) 2195. (20)
Noyes, Jr., W. A. J . Am. Chem. SOC.1932,54, 4143. Gedye, G.; Rideal, E. J . Chem. SOC.1932, 135, 1160. Welge, H.; Beckman, A. J . Am. Chem. SOC.1936, 58, 2462. Melville, H. W. Proc. R . Sac. London, A 1940, 175, 164. McDonald, C. C.; Gunning, H. E. J . Chem. Phys. 1955, 23, 532. Takamuku, S.; Back, R. A. Can. J . Chem. 1964, 42, 1426. Hoffman, M.; Goldwasser, M.; Damour, P. J . Chem. Phys. 1967, 47,
Gordon, M. S. Chem. Phys. Lett. 1986, 126, 451. (21) Luke, B. T.; Pople, J. A,; Krogh-Jespersen, M. B.; Apeloig, Y.; Chandrasekhar, J.; Schleyer, P. v. R. J . Am. Chem. Soc. 1986, 108, 260. (22) Magnusson, E. Aust. J . Chem. 1986, 39, 735. (23) Fronzoni, G.; Galasso, V. Chem. Phys. 1986, 103, 29. (24) Glidewell, C.; Thomson, C. J . Comput. Chem. 1983, 4, 9. (25) Hopkinson, A. C.; Lien, M. H. J . Mol. Struct. 1983, 92, 153. (26) Hendewerk, M. L.; Frey, R.; Dixon, D. A. J . Phys. Chem. 1983, 87, 2026. (27) Stock, A.; Somieski, C. Eer. 1921, 5 4 8 , 740.
0 1987 American Chemical Society
