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J. Phys. Chem. 1985.89.4903-4905
spectra of alkyl olefins are more complicated than previously acknowledged, though the complications may be evident only in certain highly alkylated systems. One tempting clue as to the identity of the second valence excitation in BCH comes from the photoelectron work of Heilbronner et al. on the "ribbon orbitals" of species containing the cyclohexane ring." The orbitals in question are related to the highest-filled M O s of cyclohexane (4eJ and are composed of n(CH,) and o(CC) local orbitals! According to the photoelectron study. the ribbon M O s are separated from the tangle of the other T , o M O s by an appreciable gap in compounds containing the (14) S. F. Mason. 'Molecule Optical Activity and the Chiral Discriminations". Cambridge University Press. Cambridge. 1982. (15) E. Hcilbronncr. E. Honcggcr. W . Zambaeh. P. Sehmitt. and H. COnlhcr. Hdn Chim. ACII?.61. 1681 (1984).
cyclohexyl group. The ribbon M O s of BCH are observed at 9.8 eV in the photoelectron spectrum of Figure 1. As a result of their low ionization potential, the ribbon orbitals "are good donator orbitals and ... they must play an important, if not dominant role in intramolecular charge transfer processes".Is Extension of this idea leads directly to the assignment of the two low-lying valence T * , the latter k i n g transitions of BCH as r T* and n(CH,) largely an intramolecular charge-transfer transition from the cyclohexyl group to the double bond. Both of these transitions are allowed with long-axis polarization.
-
-
Acknowledgment. Financial support of the Swiss National Science Foundation (Project No. 2.219-0.84) and the United States National Science Foundation (Grant CHE-8416312) are gratefully acknowledged, as are sweral discussions of these spectra with Dr. Aharon Gedanken.
Shock-lube Study of the Rate Constant for Excited OH"(,X+) Formation in the N,O-H, Reaction Yoshiaki Hidaka,* Hirokazu Takuma,' and Masao Sup' Advanced Instrumentation Center for Chemical Analysis. and Department of Chemistry. Ehime Uniuersiry. Bunkyo-cho. Maisuyama 790. Japan
The OH' zZ+-znchemiluminescence in N , G H , highly diluted with Ar was studied over the temperature range 140W2WO K behind reflected shock waves. From the relationship between OH' concentration and temperature in three mixtures it was determined that OH' was mainly formed by H + 0 + M = OH' + M ( I ) and N,O + H = N, + O H * (6). The rate constant of reaction 6 was found to be (1.6 I ) X IO" exp(-50300/Rn cm' mol-' s-'.
*
Introduction The mechanism and rate constant expressions for electronically excited OH'(zX+) formation in the high-temperature oxidation of H, by 0,have been studied in detail.ld It was found to be formed mainly by the reaction H + 0 + M = O H * + M with k = I .2 X IO1' exp(-6940/R7) cm6 mol-z s&." The OH' emission in the high-temperature oxidation of H, by N,O has k e n studied only by Soloukhin,' who inferred H + N,O = OH* + N, a s the reaction forming OH' and discussed the relationship between OH* and O H . A detailed study of the rate constant of O H * formation in the high-temperature oxidation of H, by N,O has not yet been reported. This work, therefore, was carried out to clarify the identity and the rate constant of the reaction forming OH' in the hightemperature oxidation of H, by N,O. Experimental Section The apparatus and procedures used in this study are essentially the same as described previously! Emission was monitored with a Hamamatsu R-306 photomultiplier after passing through an interference filter (Amax = 3055 A, half-width = 150 A) located behind a quartz window o f t h e shock tube and through two slits either 0.18 and 0.24 mm wide separated by a cylindrical tube 2 cm long. Infrared emission through two 0.8-mm slits and an = 4.68 r m , half-width = 0.1 r m ) was interference filter (A,, observed with Fujitsu IV-2OOC4 InSb detector to learn the time variation of N,O concentration. The ultraviolet and infrared emissions were observed simultaneously behind reflected shock waves. The O H a concentration was determined from the OH' emission intensity by using the relationship between O H * concentration and voltage found before." Data interpretation was carried out by computer simulations essentially as described previously! 'Advanced Instrumentation Center for Chemical Analysis. 'Department of Chemistry.
0022-3654/85/2089-4903$01.50/0
Figure 1. Typical oscillograms a l N,O IR thermal emission (upper trace) and O H * U V chemilumincrcence cmission (lower tract). The sweep speed i s 100 Irr/division: mixture B. I; = 1612 K. P, = 2.14 atm.
Three reaction mixtures were employed: (A) 2.0% N,O. 1.0% H,. 97.0% Ar: (B) 1.0% N,O, 1.O% H,. 98.0%: Ar: (C) 1.0%N,O. 0.5;H,. 98.5% Ar. The H, and Ar;specified to.be 99.9% and 99.99% pure, respectively, were obtained from commercial cylinders and used without further purification. The commercial( I ) W. C. Gardiner. Jr., K. Morinaga. D. L. Riplcy, and T. Takcyarna. Phys. Fluids Suppl., I, 120 (1969). (2) D. Gutman. R. W. Lutr. N . F. Jacobs. E. A. Hardwidge. and G. L. Schou. J. Chcm. Phy.v.. 48. 5689 (1968). (3) S. Ticktin. C. Spindler. and H. I . Schiff, Discuss. Foroday Sm..44. 218 (1961). (4) T. Koike and K. Morinaga. Bull. Chem. SOC.Jpn., 49. 1457 (1916). ( 5 ) T. Koikc and K. Morinaga. Bull. Chem. Soe. Jpn., 55. 52 (1982). (6) Y.Hidaka. S. Takahashi. H. Kawano. M. Suga. and W. C. Gardiner. Jr., J . Phyr. Chew., 86, 1429 (1982). (7) R. L. Soloukhin, Symp. ( l n l . ) Combust.. [PPx.]. 141h. 77 (1913).
0 1985 American Chemical Society
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The Journal of Physical Chemistry, Vol. 89, No. 23, 1985
Letters
t
v)
..
1
E w
101
Figure 2. Observed maximum [OH*] and computed maximum [OH*] vs. reciprocal temperature: (0)mixture A; (0)mixture B; (A) mixture C; (-) calculated with k, = 1.1 X lo', exp(-50300/RT) cm3 mol-' s-l and literature values for k2, k,, k4, and k,; (--) calculated with k6 = 1.6 X lOI4 exp(-50300/RT) cm3 mol-' s-l and revised rate constant expressions for k,, k,, and k,; (-.-) calculated without reaction 6.
1000
1
1
I
I
I
I
Figure 4. Comparison of observed and calculated ignition delay times t,. The calculation was carried out with revised expressions for k,, k,, k,, and k6: (0)mixture A; (0)mixture B; (A)mixture C; (-) calculated for mixture A; (--) calculated for mixture B; (-.-) calculated for mixture C.
.
3
I
0
F r
U 0
O O
200
400 t/ys
M I
(b)
V E 4
0 E W
I
I
I
I
Figure 3. Comparison of the observed and calculated ignition delay times (0)mixture A; (0) mixture 8;(A)mixture C; (-) calculated for mixture A; (--) calculated for mixture B; (- * -) calculated for mixture C.
rm using literature values for k2, k,, k,, and k,:
grade N,O was purified by trap-to-trap distillations, and a purity of 99.9% or higher was confirmed by a gas-chromatographic analysis. Results and Discussion A typical oscillogram is shown in Figure 1. At low temperatures, the OH* emission intensity behind reflected shock waves increases slowly and then decreases slowly. At high temperatures, the emission intensity rose rapidly to a maximum and fell first quickly and then slowly with a decay rate similar to that shown in Figure 1. The maximum OH* concentration was calculated from the OH* maximum emission signal. The relationship be-
t/ys Figure 5. Comparison of [OH*] profiles computed by using revised expressions for k,, k4. k,, and k, with observed emission profiles. (a) Mixture B, T, = 1684 K, P , = 2.29 atm: (-) observed; (- - -) calculated. (b) Mixture A, T , = 1394 K, P , = 1.76 atm: (-) observed; ( - - - ) calculated.
tween maximum OH* concentration and reciprocal temperature is shown in Figure 2. The electronically excited OH* in H2 oxidation by O2is formed mainly by the reaction
H
+ 0 + M = OH* + M
(1)
below 2000 K.6 In H2 oxidation by NzO, reaction 1 is also responsible for some OH* formation, because H and 0 also exist there. When the H2-02 mechanism6 is augmented to include the N20-H2 reactions
J . Phys. Chem. 1985, 89, 4905-4908
+ M = N2 + 0 + M N 2 0 + 0 = N2 + 0 2
NzO
+ 0 = NO + N O N20 + H = N, + OH
N20
from our study of N 2 0 decomposition and the N20-H2 reaction? the calculated t , values are in good agreement with the observed ones, as shown in Figure 4. The revised rate constant expressions used are k3 = 7.0 X l O I 4 exp(-28000/RT) cm3 mol-’ s-l, k, = 5.6 X l O I 4 exp(-28000/RT) cm3 mol-’ S-I, and kS = 1.5 X l O I 4 exp(-15000/RT) cm3 mo1-l s-l. When these values are used, the rate constant calculated for k6 becomes 1.6 X l o i 4 exp (-50300/RT) cm3 mol-’ s-l. Both the OH* concentration and ignition delay time t , calculated with these values fit the experimental data well as shown in Figures 2 and 4. Systematic errors in the instrumentation and kinetic model (including the rate constant of the quenching reactions) influence the k6 value. The rate constant expression for k6, therefore, may be assigned generous error bounds as (1.6 f 1) X 1014exp(-50300/RT) cm3 mol-’ s-I. A comparison of observed with computed OH* emission profiles using the mechanism including the revised rate constants for k3, k4, kS,and k6 is shown in Figure 5. The calculated and observed curves are in fairly good agreement. Since our k6 expression is derived from fitting the intensity maxima only (Figure 2), the agreement with the time dependence (Figures 4 and 5) provides an independent confirmation of k6 and also of the predictive ability of the mechanism and rate constant expressions used.
(2) (3) (4)
(5) with literature values for their rate constantss, the maximum OH* concentrations calculated at high temperatures are much smaller than observed, as shown in Figure 2. Therefore, a large portion of OH* is formed by some other reactions. The simplest assumption is N 2 0 H = OH* N 2 A”,= 17 kcal/mol (6)
+
4905
+
Though reaction 6 is endothermic, it might be responsible for production of OH*. In computing OH* concentrations, it is necessary to assume the quenching rate constant of OH* by each species. The values reported previously6were used, and we assumed that N O and N2 have the same quenching rate as O2 and that N 2 0 has same s quenching rate as H 2 0 . We adopted the value of 0.7 X reported previously6 as the radiative lifetime of OH*. When the value of k6 = 1.1 X 1014exp(-50300/RT) cm3 mol-I s-l is used as the rate constant of reaction 6, the OH* concentrations calculated fit the experimental data best, as shown in Figure 2. Ignition delay times t , were defined as times between reflected shock arrival and maximum OH* emission intensity. A plot of t , vs. 1 / T for each mixture is shown in Figure 3. The t , values calculated with the literature rate constants are longer than observed. This is considered to come from a real shortcoming of the rate constant expressions in the N20-H2 reaction. When we use revised rate constant expressions for k,, k4, and k5 obtained
Conclusion Chemiluminescent OH* formation in H2 oxidation by N 2 0 over the temperature range 1400-2000 K was found to be mainly due to H 0 + M = OH* M and N 2 0 + H = N 2 + OH*. The expression k6 = (1.6 f 1) X lOI4 exp(-50300/RT) cm3 mol-’ was determined for reaction 6. With this rate constant, the OH* emission intensities and ignition delay times were satisfactorily interpreted.
+
(8) D. L. Baulch, D. D. Drysdale, D. G. Horne, and A. C. Lloyd, “Evaluated Kinetic Data for High Temperature Reactions”, Vol. 2, Butterworths, London, 1973.
+
(9) Y. Hidaka, H. Takuma, and M. Suga, Bull. Chem. SOC.Jpn., 58,291 1 (1985).
Gas-Phase Oxidation of Silver: The Reaction of Silver Clusters with Ozone James L. Gole,* R. Woodward, J. S. Hayden, High Temperature Laboratory, Center for Atomic and Molecular Science, and School of Physics, Georgia Institute of Technology, Atlanta, Georgia 30332
and David A. Dixon Central Research and Development Department, E . I. DuPont de Nemours and Company, Wilmington, Delaware 19898 (Received: June 26, 1985) In a device which produces a metallic flow intermediate to that of a standard effusive source (atoms and small percent diatomics) and laser vaporization-plasma formation followed by rare gas entrainment at pressures exceeding several torr (wide cluster distribution), the reaction of small silver clusters, M,, n 2 3, with ozone has been observed in the gas phase at pressures in the millitorr range. Chemiluminescent emission from the products Ago and what appear to be Ag20 and Ag,O, where x 2 3, has been monitored. We associate a red shift of the spectral features with the increasing size of the metal cluster oxide. The spectral data combined with supplementary thermodynamic information demonstrate that Ago* with enough energy to account for the observed chemiluminescence cannot be produced through the reaction of either Ag or Ag, with 0 3 .The smallest cluster whose reaction can yield excited states of A g o is the trimer. The formation of Ag20* can also be achieved through reaction of the trimer; however, it may also be accounted for via reaction of higher clusters. There has been widespread and growing effort to understand the structures and properties of small free metallic clusters. An increasing number of experimental characterizations1 are now beginning to balance the impressive array of theories which have been applied to these systems.2 Small clusters have been generated in flow systems, reacting with reagents in another continuous or pulsed flow stream under high pressure (-30-500 torr) conditions in a modified merged flow e n ~ i r o n m e n t . ~The products in the flow have been measured mass spectrometrically; however, this technique provides no direct measurement of structural or dynamic properties. In contrast, we have recently detected chemiluminescence from the oxidation of sodium clusters by halogens under single-collision condition^.^ This latter experiment is ad0022-3654/85/2089-4905$01.50/0
vantageous in that it provides both dynamic and spectroscopic information about the products of the metathesis. The chemistry of silver is of substantial technological importance in both photographic5and catalytic processes.6 This is exemplified through the use of supported silver in the epoxidation of ethylene and bulk silver in the dehydrogenation of methanol to give formaldehyde. The bulk of experimental studies, thus far, on “naked” silver clusters have been done in low-temperature matrices.’,* We report here initial results from a study of the oxidation of gas-phase silver clusters with ozone. The present approach provides a novel route to the preparation of small metal cluster oxides of silver, Ag,O. The apparatus used in this study is a modified version of that 0 1985 American Chemical Society
