2810
The Journal of Physical Chemistty, Vol. 83, No. 22, 1979
Michael et ai.
Rate Constant for the Reaction H 4- NOp from 195 to 400 K with FP-RF and DF-RF Techniques J. V. Michael,+ D. F. Nava, W. A. Payne, J. H. Lee,$ and L. J. Stlef" Astrochemistry Branch, Laboratory for Extraterrestrial Physics, NASA/Goddard Space F//ghtCenter, Greenbelt, Maryland 2077 1 (Received May 23, 1979) Publication costs assisted by NASA
Absolute rate constants for the reaction of atomic hydrogen with nitrogen dioxide have been measured over the temperature range 195-400 K with two independent techniques both of which utilize resonance fluorescence detection of H.The reaction was studied from 230 to 400 K with the flash photolysis-resonancefluorescence (FP-RF) technique. No variation with temperature was observed, the best representation for the rate constant cm3molecule-' s-'. The discharge flow-resonancefluorescence (DF-RF)technique being (1.32 f 0.12) X was also used to study the reaction over the temperature range 195-368 K. The results are best represented by the temperature-independent value of (1.50 f 0.26) X lo-', cm3molecule-' s-l. The mean value of the two determinations for 195 IT I400 K is then k1 = (1.41 f 0.26) X cm3 molecule-' s-' where the error is two standard deviations. The combined result is compared with previous results, both direct and indirect. The reaction is also considered theoretically,especially with regard to the question of the temperature dependence and the absolute magnitude of the rate constant.
Introduction The reaction between hydrogen atoms and nitrogen dioxide H + NO2 --.+ OH + NO (1) with AHomS= -29 kcal/mol has drawn increased attention for several important reasons. Studies of the absolute reaction rate constant14 and the energy states of the products of the reaction7-12are of particular interest since it is now known that vibrationally excited OH is produced in reaction L7-lo In addition to this fundamental interest, reaction 1 has been used as a source for vibrationally excited OH for chemical kinetic studies of that species itself.1° Reaction 1 has also been used in discharge flow studies to prepare ground-state OH,13 and many of the currently accepted rate constants of OH with molecules and radicals come from this type of study.14 Lastly, reaction 1 has served as the major calibration reaction for H atoms in discharge flow experiments. Secondary processes become important if the titration is carried out at relatively high [HI,, and then the overall stoichiometry changes from 1:l to 1:1.5 H atoms to NO2 molecules. These secondary complications have been previously discussed.16 In this work we have measured the rate constant for reaction 1 over significant temperature ranges with the flash photolysis-resonance fluorescence (FP-RF) technique and also with the discharge flow-resonance fluorescence (DF-RF) technique, Our results can be compared to six earlier studies,'* four of which are Reinterpreted results4i6from the relative experiments of Ashmore and Tyler' give k l = 2.0 X lo-', cm3molecule-l s-l at T = 633 K. Rosser and Wise's relative results? also reinter~reted?~ yield 1.6 X lo-', cm3molecule-' s-l at T = 500 K. The first direct study was by Phillips and Schiff3 who used the discharge flow-mass spectrometric technique. They report 'Visiting Professor of Chemistry, Catholic University of America (CUA), Washington, DC 20064. t Research Associate, CUA, Washington, DC 20064; currently with the Department of Energy and the Environment, Brmkhaven National Laboratory, Upton, Long Island, NY 11973. *Adjunct Professor of Chemistry, CUA, Washington, DC 20064. This article not subject to U.S. Copyright.
kl = (4.8 f 1.2) X cm3 molecule-l s-l at 298 K. Bemand and Clyne6with the DF-RF technique report kl = (1.1 f 0.3) X lo-', cm3 molecule-l s-'. Clyne and Monkhouses and Wagner, Welzbacher, and Zellner4 measured the temperature dependence of kl with the same technique (DF-RF). Their results agreed closely as shown in the respective Arrhenius expressions kl = (4.8 f 1.0) X exp(-4Ol f 71/T) cm3molecule-' s-' for 298 I T 5 652 K and kl = (7.1 f 3.0) X exp(-505 f 84/T) cm3 molecule-' s-l for 243 I T 5 461 K. Even though these recent direct results are in good agreement, we note that both studies were carried out with the same technique. Since it is important to study chemical reactions with more than one technique, reaction 1 was investigated with the FP-RF technique over the temperature range 230-400 K. The results, as reported here, do not agree with the earlier determinations, and we accordingly have made a separate set of measurements with the DF-RF technique from 195 to 368 K. These results are also reported here. Experimental Section Flash Photolysis-ResonanceFluorescence (FP-RF). The flash photolysis-resonance fluorescence apparatus which was used in the present study and its application to hydrogen atom reactions have been described previously in detail.16-18 Thus, only those points which are specific to the present study will be given here. Scattered fluorescent photons at Lyman a (121.6 nm) were proportional to ground-state [HI and were observed at right angles to both the resonance lamp and the photoflash lamp through 5 cm of flowing dry air. The detector was an EMR solar blind photomultiplier. In all of the reported work [HI, < l o l l atoms ~ r n - ~The . photodecomposition of CHI (A > 110 nm, LiF cutoff) served as the source of H in these experiments. Ternary mixtures of NO2, CH4,and diluent Ar were then anticipated for the reaction studies. Since NO2is a fairly reactive molecule, it was necessary to establish that it was stable in the system for a sufficient length of time so that no ambiguity could arise regarding concentration measurements. Therefore, the system was tested with purified NO2, and no measurable pressure loss (MKS Baratron capacitance manometer) with time was observed. Also
Published 1979 by the American Chemical Society
Rate Constant for the Reaction H iNO2
absorption coefficients at 313.0 and 435.8 nm were measured at T = 298 K over a path length of 154.5 cm. The light source was a low-pressure Hg lamp with corresponding interference filters. We obtained 6 3 1 3 . b = 5.16 f 0.06 cm-' (atm at 273 K)-l, base e, and ~ 4 3 5 . 8=~17.8 ~ f 0.3 cm-l (atm at 273 K)-I, base e . The 313.0-nm value compares favorably with Bass et al.19 and Johnston and Graham,2O being within 6% of these determinations. The small differences are probably due to absolute pressure accuracy. The most recent published measurement of t435.8nm is due to Hall and Blacet21 whose results are uniformly higher at all wavelengths than those of Bass et al.19 by 10-20%. Hall and Blacet's graph interpolates to e435.8nm = 19.5 cm-I (atm at 273 K)-l, base e. This value is higher than the present value by l o % , and, therefore, the present value a t 435.8 nm is consistent with findings of Bass et al. The spectroscopic experiments clearly showed that NO2 was stable. We then determined, again by absorption at 435.8 nm, that NOZ-Ar-CH4 high-pressure mixtures were stable for periods of days in the blackened 25-L storage bulbs. Finally, with an N02-Ar mixture of mole fraction XNoB= 2.50 X the minimum flow rate through the stainless steel tubing and brass reaction cell was established so that no NO2was heterogeneously lost. [NOz]was measured before and after passage through the cell again by absorption at 435.8 nm. The experiments to establish minimum flows were carried out as a function of temperature. We found a substantial loss of NO2 if the temperature was greater than 420 K. Since the rate constant for reaction 1 is so large, small concentrations of NO2 had to be accurately measured for the reaction studies. We elected to use the premixed, XNo2 = 2.50 X N02-Ar mixture to prepare the reaction mixture by dilution techniques. The 435.8-nm absorption of the premixture was always checked prior to reaction mixture preparation. Accurately measured amounts of this mixture along with measured amounts of CH4and diluent Ar then served as the reaction mixtures. Unfortunately, the levels of NOz in these diluted mixtures was too low to measure spectrophotometrically under reaction conditions; however, experience with the more concentrated NO2CH4-Ar mixtures gave confidence khat no NO2 would be lost and the mixture composition would be stable in storage. Furthermore, if the flow rates of the mixture were maintained well above the minimum levels established with the more concentrated mixture, the composition of the diluted mixtures would remain constant and be as calculated. Note that the NO2 concentration calculation requires five separate pressure measurements: two for the premixture, two for the diluted mixture, and one for the total pressure under reaction conditions. All pressure measurements were made with an MKS Baratron capacitance manometer. According to the manufacturer, each reading is absolutely accurate to f0.5%, and, since five readings are involved, the uncertainly in [NO,] is estimated to be k3%. Note further that under no circumstances was P N Ohigh ~ enough so that corrections became significant due to the equilibrium 2NOZ P N204.22 Preliminary reaction studies showed that the observed first-order decay constants, obtained in the usual way from the exponential decay of counts with time, were independent of substantial variations in photoflash intensity (Le., [HI,) and flow rate of the mixture through the reaction cell. As in earlier studies, exactly identical experiments with no added NOz were performed under each experimental condition in order to measure the diffusional loss contribution of H atoms to the observed decay con-
The Journal of Physical Chemistry, Vol. 83, No. 22, 1979 2819
stants when NO, was present. Measurements of kl were performed at five temperatures ranging from 230 to 400 K, the upper temperature being dictated by heterogeneous loss of NO, in the reaction cell. Discharge Flow-Resonance Fluorescence (DF-RF). The DF-RF apparatus and methods have also been described previously.23Application to H atom reactions has also been r e p ~ r t e d ,so~ only ~ , ~those ~ details specific to the present study will be given. H atoms were observed by Lyman-a fluorescence through 5 cm of flowing dry air by an EMR solar blind photomultiplier. 1.00% mixtures of NO2 in He served as the source of reactant. Since reaction 1 is a well-known titration reaction,15routine H atom calibration experiments were performed for each experimental condition. Typical sensitivities for atom detection ranged from 0.5 X lo9 to 2 x io9 atoms ~ r n at - ~ a signal-to-noise ratio of unity, depending on the cleanliness of the optical windows. The kinetic experiments were performed from 195 to 368 K at pressures from 1.7 to 3.2 torr, the carrier gas being He. Temperature variation was accomplished with cooling or heating fluids contained in a polyurethane trough surrounding the flow reactor. Decay constants were measured over atom concentrations ranging typically between 1 X 1O1O and 2 X loll atoms ~ m - ~ Extrapolation , of decay constant results to zero time gave for each experiment an estimate of [HI,. This estimate and the measured [NO,], allowed the ratio [NO,],/[H], to be calculated. Materials. Argon (Matheson, 99.9995%),helium (Airco, 99.9999%), and methane (Matheson, 99.97%) were used without further purification. One hundred torr of NOz (Air Products, 99.5%) was combined with 660 torr of research grade Oz (Matheson, 99.99%) for at least 12 h at room temperature. The condensible NO2 was frozen out at liquid N2temperature and the O2was pumped away. NO2 was further purified by bulb-to-bulb distillation, with the middle third being retained. The pure white solid was stored in a thoroughly blackened bulb.
Results FP-RF System. In all experiments [NO,], >> [HI, so that first-order kinetics are applicable for H atom decay. Since counts are proportional to [HI, the relative fluorescent signal decay is a measure of relative [HI. Then In [HI = -kohsdt+ In [HIo (2) where kobsd = k1[N0210 + kd (3) kl is the bimolecular rate constant for reaction 1 and kd is the rate constant for diffusional loss of H atoms out of the intersection volume. Figure 1presents typical first-order decay results at 400 K. Similar results with [NOz], = 0 under identical conditions then give kd as shown in eq 3. All decay constants are evaluated by linear least-squares methods, A knowledge of kobsd and kd then allows kl to be calculated. Repeat determinations under a given experimental condition, but with photoflash intensity variation, then give an estimate of k1 along with a standard deviation for that condition. The results at five temperatures from 230 to 400 K are given in Table I. As seen in the table, no variation of kl with temperature was observed. The best representation is kl = (1.32 f 0.12) X lo-', cm3molecule-l s-l for 230 IT I400 K where the error is the average of the one standard deviation values shown in the table. DF-RF System. The absolute sensitivity for H atoms was initially measured by titration with NOz for each experimental condition so that the pseudo-first-order
2820
The Journal of Physical Chemistty, Vol. 83,No. 22, 1979
Michael et al.
TABLE I: Rate Data for t h e Flash Photolysis-Resonance Fluorescence Study of the Reaction H t NO, T,K
[NO,], mtorr [CH,], mtorr
230
263
298
343
a
flash energy, J
0.075 0.15 0.11 0.23
120 240 180 360
30 60 90 90
95-203 81-203 127 95-163
0.13 0.19 0.28 0.38
160 240 360 480
40 60 90 120
36,81 36,81 36,81 81
0.09 0.31 0.19 0.28 0.38
120 200 240 360 480
30 50 60 90 120
81-182 81,163 68-127 127-182 81-163
0.21 0.31 0.41 0.46
400
[Ar], torr
40 60 80 90
160 240 320 360
0.14 0.21 0.28 0.21 0.31 0.41 0.55 0.31 0.46 0.62 0.41 0.62 0.52
4
2
2 3 -2 9 6 4 3 3
4
60
90
90 90 120 120 150
1.38f 0.11 1.31 f 0.00 1.23i 0.09 1.24i 0.12
20 2 2 2 2
81 81 36,81 28,81 81 28-81 15-68 81 36,81 36,81 36,81 81 36,81
60 60 80
1.32f 0.08 1.30i 0.10 1.52 1.24f 0.15 1.31 f 0.1lb 1.51i 0.08 1.50f 0.18 1.30 i 0.11 1.28f 0.13 1.38f 0.15b 1.41i 0.08 1.44* 0.04 1.27f 0.13 1.46f 0.08
4 1 3 12
36,81 36,81 36,81 36,81
40 40 40
160 160 160 240 240 240 320 360 360 360 480 480 600
k,: cm3 molecule-' s-'
no. of expt
8
+%$-%
1 1 2 2 1 3 3 1 2 2 2 2 2 24
1.30 1.57 1.42f 0.06 1.12 i 0.02 1.47 1.37i 0.06 1.30f 0.09 1.12 1.39i 0.17 1.31f 0.08 1.20i 0.14 1.36i 0.05 1.12i 0.08b 1.30 i 0.15
Average k , a t that temperature.
Error limit is the standard deviation.
r------
\
22 0 3 / 1 1 " " 0
I
2
3
4
5
6
t /ms
Figure 1. First-order decay plots with the FP-RF apparatus at 400 K: ( 0 )PT= 40 torr, PN02= 0.14 mtorr, PCH,= 160 mtorr; kobsd= 575 f 16 si; (0) pT= 60 torr, f N O = 0.21 mtorr, f,, = 240 mtorr; k&d = 656 19 s-1: (A) = 9d torr, P,, = 0.31 mtorr, pCH= 360 mtorr; kobsd= 912 f 26 s-l: (0) PT= 750 torr, PNOI= 0.51 mtorr, P,,, = 600 mtorr; k&, = 1366 rt 25 s-'. Lines are from linear least-squares analysis.
*
relation, [NO,], >> [HI,,could be maintained in the kinetic runs. Under these conditions the decay of atoms is then given by In [HI= -kobsdd/u + In [HI, (4)
I
I
10
20
1
30 d /crn
1
1
40
50
Flgure 2. First-order decay plots, In [HI against d, with DFRF apparatus ~ , = 123 f at 368 K: (0)[NOz] = 7.08 X 10" molecules ~ m - kobd 3 s-'; (A)[NO,] = 1.02 X lo'* molecules om3, k, = 188 f 3 s-l: ( 0 )[NOz] = 1.33X 10" molecules ~ m - khb, ~ , = 247 f 3 Si; (A) [NO,] = 1.68 X IO1*molecules ~ m - k,~ , = 310 f 6 s-'. Lines are from linear least-squares analysis. PT= 2.04torr, v = 285 1 cm S-1.
where d is the distance from the probe to the detector and u is the linear flow velocity. Figure 2 shows some typical decay plots at 368 K where the lines are determined by linear least-squaresanalysis. The observed decay constants are then corrected for axial diffusion in the usual way (Le.,
Rate Constant for the Reaction H
+ NOp
The Journal of Physical Chemistry, Vol. 83, No. 22, 1979 2821
TABLE 11: Rate Data for the Flow Discharge-Resonance Fluorescence Study of the Reaction H t NO, [NO,], molecules [NO,], / T, K P,torr v , cm s-l [HI, kcor? s-l slopeb and interceptC from eq 5 6.46 6 88 f 2 S = (1.242 0.08)X lo-"? I = 17 f 10 195 1.69 1459 8.59 7 129 f 3 10.3 8 149 i 2 13.5 9 187 i 11 17.9 13 233 i 2 6.75 5 88 ?: 1 S = (1.61i 0.07)X I = -24 + 11 985 3.00 9.00 5 130 f 4 7 127 f 2 9.93 5 170 i 5 11.2 155 ?: 5 12.6 9 13.9 6 203 i 5 15.3 7 222 f 10 18.5 7 279 f 3 18.5 8 258 f 5 21.2 6 332 f 7 24.1 6 363 f 9 Results at 195 K, k , = (1.4i 0.3)x 7 130 f 6 S = (1.59f 0.05)X lo-"? I = -2 f 12 2186 298 1.82 7.98 5 126 f 3 8.33 10.1 6 149 f 1 11.0 6 199 f 8 12 286 f 4 19.1 21.2 9 312 i 4 23.4 19 400 f 38 15 375 f 9 24.2 13 499 f 2 32.6 17 602 f 12 37.4 25 633 f 41 39.4 4 1163 80 f 3 S = (1.49c 0.10)X 10-"? I = -7 f 13 3.16 6.05 5 116 f 1 7.94 9.58 6 129 f 3 11.8 5 172 f 5 4 186 f 2 13.4 5 277 f 3 17.7 20.5 3 283 f 2 Results at 298 K, k , = (1.5i 0.1)x 368 2.04 2851 5.90 6 107 f 3 S = (1.67i 0.16)X 10-lo,u I = 11 i 20 7 .08 8 126 i 3 10 157 f 2 8.60 9 146 f 9 9.68 13 10.2 194 i 3 12 12.4 202 i 5 10 12.5 237 f 3 13 13.3 258 f 3 17 233 i 10 14.6 16 327 i: 6 16.8 13 342 i 6 21.2 1447 2.17 6.25 9 91 i 2 S = (1.41i 0.07)X 10-"? I = -4 f 14 13 94 f 3 7.92 10 137 f 1 9.80 11.4 19 177 f 1 13.3 11 166f 2 15.7 16 199 f 3 16.1 12 213 f 1 17.0 11 257 .I 5 19.1 16 236f 6 21.4 16 339 f 4 24.0 18 338 f 4 28.0 12 361 f 4 31 .O 19 424 f 18 32.0 13 463 f 6 Results at 368 K, k , = (1.5 f 0.2) X 10-lOa a Error limits are one standard deviation values. Slope S = h , in units cm3 molecule-Ls-,. Intercept Z = the rate constant for wall termination k , in units s - l .
kc,, = kobsd(1 + k&lldD/u2)26with the known diffusion coefficients26and the experimentally measured values of kom and u. These corrections were typically 5-15%; however, for the highest temperature, lowest flow velocity, and largest [NO,], the corrections were -25%. The corrected decay constants at each experimental condition
are given in Table 11. Also given in the table are the linear least-squares analyses of the data according to the equation (5) kco, = kiD"zlo + k, k, is the rate constant for wall termination, and kl is the rate constant for reaction 1. Note that the slopes from the
2822
The Journal of Physical Chemistry, Vol. 83, No. 22, 1979
linear least-squares analysis are the values of kl while the intercepts correlate with k,. Inspection of the results in Table I1 shows that there is no variation of kl with temperature between 195 and 368 K within experimental error. The best representation for this temperature range is kl = (1.50 f 0.26) X cm3 molecule-' s-' where the error is chosen to overlap all of the determinations at any temperature.
Discussion The present FP-RF result of k1 = (1.32 f 0.12) X 10-lo cm3 molecule-' 5-l for 230 IT I400 K and the DF-RF result of kl = (1.50 f 0.26) X cm3 molecule-* s-' for 195 IT I368 K are in good agreement. Both techniques show that reaction 1 has no significant temperature dependence. The simple average of the two results for the present temperature range of 195 IT I 4 0 0 K is then kl = (1.41 f 0.26) X cm3 molecule-' s-l where the error is taken at the two standard deviation level. This combined result can be compared to the previous results of both relative and direct experiments.14 Ashmore and Tyler's' and Rosser and Wise's2 reinterpreted relative results4s5are 2.0 X (T = 633 K) and 1.6 X (T = 500 K), respectively, both in units of cm3 molecule-' s-l. The direct studies have all been carried out with the flow reactor technique. Phillips and Schiff with mass spectrometric detection report (4.8 f 1.2) X lo-'' cm3 molecule-' s-' (T = 298 K).3 The room temperature values of Clyne and Monkhouse5 and Wagner, Welzbacher, and Zellner4 are, respectively, (1.32 f 0.18) X and (1.28 f 0.25) X lO-'O cm3 molecule-l 8. Thus, the present results are in good to adequate agreement with the relative results and with all of the room temperature direct results except those of Phillips and Schiff. The present temperature independence of kl,however, disagrees with the higher temperature direct values of both Clyne and Monkhouse5and Wagner, Welzbacher, and Zellner4whose results suggest a definite temperature dependence. They report respective Arrhenius expressions of (4.8 f 1.0) X exp(-401 f 71/T) and (7.1 f 3.0) X exp(-505 f 84/T) cm3molecule-' s-'. This disagreement is puzzling particulary since both the Clyne and Monkhouse and Wagner, Welzbacher, and Zellner data were obtained by the DF-RF technique which is one of the present techniques. Comparison of their experiments with the present also indicates similar conditions of total pressure, [NO& and [HI,. We do note that the major discrepancy between Clyne and Monkhouse's results and those reported here appears to be only at their two highest temperatures. The question of temperature dependence for reaction 1 is of theoretical significance as recently discussed by Kaufman and LevineaZ9These authors suggest that the disequilibrium product energy states can be used to estimate the decrease in entropy of a reaction, and hence the decrease in the A factor, from that expected for a statistical distribution. They state that the nearly statistically distributed energy dissipation in reaction 1 and the high A factor from Clyne and Monkhouse5 and Wagner, Welzbacher, and Zellner4 are consistent with the nonstatistical energy dissipation (0.90,0.03, and 0.07 fractions of the available energy in vibrations of OH, rotations of OH, and relative translations plus internal excitation of Oz, respectively) and low A factor for the apparently similar reaction H + O3 OH + Oz (6) The rate constant for reaction 6 has been reported by us30 to be k g = (1.33 f 0.32) x W0exp(-449 f 58/T) cm3 molecule-l s-'. This has been corroborated by Keyser3' who
-
Michael et al.
reports k6 = (1.50 f 0.18) X exp(-499 f 32/T) cm3 molecule-' s-l. Neither of these studies are in agreement with the earlier work of Clyne and Monkh~use.~ The energy states of the products from reaction 1 have been considered from conventional crossed molecular beam analysis1'J2 and also from IR chemilumine~cence~~~ and laser fluorescences sampling of hydroxyl states. The average fractions of the available energy ending up in translations, rotation of OH, and vibrations of OH are -0.24, -0.2, and -0.4, respectively. The remaining energy apparently populates internal states of NO. A subject of particular interest has been whether the intermediate, HONOt, is tightly bound leading to snarled trajectories and subsequently to energy redistribution, more or less statistically, into all available degrees of freedom, or whether the interaction is direct. A consideration of the energy partitioning has led Silver et aLEbto prefer the direct interaction, and, by analogy with classical trajectory models,32they suggest that the potential energy surface is highly attractive with some degree of secondary encounters. The reactants and products correlate with both 'A' and 3A' electronic states with the singlet, if it is the ground state, exhibiting distinct cis and trans "wells" in the potential surface. The data and analogies with the Li + NOz reaction led Silver et al.g to seriously consider that reaction 1may proceed through the triplet state even though this suggestion disagrees with chemical intuition. Also the triplet interaction is expected to have a significant barrier. The present result shows no temperature dependence for kl and argues against the suggestion of a triplet intermediate. Also it suggests equal A factors for reactions 1and 6 and thereby damages the correlations of Kaufman and L e ~ i n e .We ~ ~point out that, even with the detailed state-to-state information on reaction 1, there is still ambiguity regarding the shape of the potential surface, the data being variously interpreted with both attractive7isb and repulsive potential surfa~es.~ The question of the deep well has not even been ruled out if the lifetime for dissociation of the intermediate is short compared to that for intramolecular energy transfer. It therefore appears to us that about all than can be said with certainty about the potential energy of interaction is that at large distances it will be dominated by the long-range potential (i.e., r4). We therefore calculate through the kinetic theory of gases the thermalized Lennard-Jones rate constant for the interaction of H with N02.27Fairly accurate estimates can be obtained if there are sufficient transport property data available so that values of a and c / k can be specified. In the present case (a, Elk) values of (2.18 A, 35.5 K) and (4.55 A, 190 K)have been used for HZ8and N02,27respectively, along with , combining rules, c12= (al + a2)/2and 6'2 = ( ~ l t 2 ) l / ~and tabulated integrals, Q*(2,2)(P)27(where T* = IZT/e12),to compute the Lennard-Jones collision rate constants for 195 IT 5 400 K. Through ZLJ= a1220*(2,2)(8akT/y)'/2 (7)33 Zu = (9.0 f 0.6) X cm3molecule-' s-' for the present temperature range. This compares to a similar estimate of 9 X cm3 molecule-' s-' for reaction 6.30 Unlike reaction 6 where electronic state correlation arguments 2A' (g ratio = l),we presume that the show zS 'Al absence of a significant barrier means that the reaction intermediate is the singlet; Le., zS + zA1 -,'A' (g ratio 1/4), With a naive estimate of the steric factor, 1 = 1, due to the symmetry of NO2, the estimated frequency factor cm3 molecule-' s-' for 195 IT 5 400 is then 2.3 X K. Since there is no energy barrier this becomes the rate
+
-
The Journal of Physical Chemistry, Vol. 83, No. 22, 1979 2823
O2 Quenching of the Rubrene Singlet State
constant estimate for reaction 1. By comparison with the observed value this estimate is only high by 60%. Though there is some ambiguity in the choice of Lennard-Jones parameters the present estimate of Zu cannot be inaccurate by more than 20%. This can be demonstrated by a similar calculation for uH = 2.93 8, and e H / k = 23.7 K which are derived from molecular beam data.34 Through cm3molecule-'s-l for 195 eq 7 , Z u = (10.6 h 0.8) X I T I 400 K which is within 16% of the estimate based on transport property data. We therefore assert that more than half of the H atoms, which impinge on the collision surface and have the necessary spin state, result in reaction. This implies a wide range of impact parameters. Through the previously cited trajectory analysis,32this result is consistent with an attractive reaction potential energy surface where the trajectories are snarled, thereby leading to substantial rotational energy release. This is substantially different from reaction 6 where similar consideration^^^ suggest a narrow range of impact parameters and almost no energy release into rotations, in agreement with chemiluminescence observations.' Acknowledgmentq J*H*L*and J*V*M* support by NASA under Grant NSG 5173 to Catholic University of America.
(8) (a) J. H. Brophy, J. A. Silver, and J. L. Kinsey, J. Chem. Phys., 82, 3820 (1975): (b) J. A. Silver, W. L. Dlmpfl, J, H. Brophy, and J. L. KinseY, bid., 65, 1811 (1976). (9) 6.K. Smith and E. R. Fisher. J. Phys. Chem.. 82. 2139 (1978). cioj J. E. Spencer and G. P. Glass, Chein. phys., is, 35 (1976). (1 1) H. Haberknd, P. Rohwer, and K. Schmidt, Chem. phys., 5,298 (1974). (12) H. Haberland, W. v. Lucadou, and P. Rohwer, Ber. Bunsenges. phys. Chem., 01, 150 (1977). (13) See, for example, J. G. Anderson, J. J. Margitan, and F. Kaufman, J . Chem. fhys., 80, 3310 (1974). (14) R. F. Hampson, Jr., and P. Garvin, Natl. Bur. Stand. Spec. Publ., No. 513 (1978). (151 . . See. for examde, A. McKenzie. M. F. R. Mulcahy, and J. R. Steven, J . Chem. Soc., Faraday Trans. 1 , 70, 549 (1974), and references cited therein. (16) R. 8. Klemm and L. J. Stief, J. Chem. Phys., 81, 4900 (1974). (17) W.A. Payne and L. J. Stlef, J. Chem. fhys., 84, 1150 (1976). (18) J. H. Lee, J. V. Michael, W. A. Payne, and L. J. Stlief, J. Chem. phys., 68, 1817 (1978). (19) A. M. Bass, A. E. Ledford, Jr., and A. H. Laufer, J . Res. Natl. Bur. Stand., Sect. A , 80, 143 (1976). (20) H. Johnston and R. Graham, Can. J . Chem., 52, 1415 (1974). (21) T. C. Hall, Jr., and F. E. Blacet, J. Chem. Phys., 20, 1745 (1952). (22) Equillbrlum constants are from data in JANAF ThermochemicalTables, 2nd ed., MU.Stand. Ref. mta Ser., Nati. Bur. Stand., No. 37 (1971). (23) J. V. Michael and J. H. Lee, Chem. fhys. Lett., 51, 303 (1977). (24) D. A. Whytock, J. V. Mlchael, W. A. Payne, and L. J. Stlef, J. Chem. f h v s . . 65. 487 (19761. (25) F. Kaufman, frog. React. Kinet., 1, 1 (1961). (26) DHHsat 1 atm is 1.35, 2.71, and 3.89 cm2 s-' for 195, 298, and 36b K. Calculations are with eq 8.2-44 of ref 27 with u , , ~ = 2.378 A and eHHe/k= 19.0 K (see also ref 28). (27) J. 0. Hlrschfelder. C. F. Curtis, and R. B. Bird, "Molecular Theory of Gases and Llquids", Wiley, New York, 1964. (28) K. P. Lynch and J. V. Michael, Int. J. Chem. Khet., 10, 233 (1978). (29) F. Kaufman and R. D. Levlne, Chem. fhys. Lett., 54, 407 (1978). (30) J. H. Lee, J. V. Michael, W. A. Payne, ard L. J. Stief, J. Chem. phys., 89, 350 (1978). (31) L. F. Keyser, J . Phys. Chem., 83, 645 (1979). (32) P. J. Kuntz, E. M. Nemeth, J. C. Polanyi, S. D. Rosner, and C. E. Young, J. Chem. fhys., 44, 1168 (1966). (33) S. C. Chan, J. T. Bryant, L. D.Spier, and B. S. Rablnovitch, J. fhys. Chem., 74, 2058 (1970). (34) R. W. Bickes, Jr., B. Lantzsch, J. P. Toennles, and K. Walaschewskl, faraday Discuss., Chem. Soc., 55, 167 (1973), and references therein. .
References and Notes P. 0. Ashmore and B. J. Tyler, Trans. Faraday Soc., 58, 1108 (1962). W. A. Rosser and H. Wlse, J. fhys. Chem., 65, 532, 2277 (1961). L. F. Phillips and H. I. Schiff, J. Chem. fhys., 37, 1233 (1962). H. Gg.Wagner, U. Welzbacher, and R. Zellner, Ber. Bunsenges. phys. Chem., 80, 1023 (1976). (5) M. A. A. Clyne and P. B. Monkhouse, J. Chem. Soc.,Faraday Trans. 2. -. 73.. 298 , ~(1977). (6) P. P. Bemanb andM. A. A. Clyne, J. Chem. Soc., faraday Trans. 2, 73, 394 (1977). (7) J. C. Polanyl and J. J. Sloan, Int. J. Chem. Klnet., Symp. 7, 51 (1975). (1) (2) (3) (4)
I
Production of Singlet Molecular Oxygen from the O2 Quenching of the Lowest Excited Singlet State of Rubrene K. C. Wu and A. M. Trozzolo' Radiation Laboratory+ and Chemistry Department, University of Notre Dame, Notre Dame, Indiana 46556 (Received August 14, 1978, Revised Manuscript Received Ju/y 30, 1979) Publication costs asslsted by the U. S. Department of Energy
The efficiencies of the rubrene-sensitized photooxidation of 1,3-diphenylisobenzofuranin benzene, acetone, cyclohexane, n-hexane, and isooctane have been measured. The measured efficiencies are analyzed with data from the oxygen quenching of the rubrene fluorescence in the same solvents. It is shown that about 1.5 molecules of singlet oxygen are produced as a result of the primary process in which a rubrene molecule in the S1state is quenched by ground state oxygen. This indicates that there is a substantial efficiency for the direct production of singlet oxygen by the oxygen quenching of rubrene S1 excited states: S1 + 302 T1 + l 0 2 . -+
Introduction There is conclusive evidence that a triplet state, "1, is formed when an organic molecule in its lowest excited singlet state, S1, is quenched by ~xygen:l-~ 'The research described herein was supported by the Officeof Basic Energy Sciences of the Department of Energy. This is Document No. NDRL-1915from the Notre Dame Radiation Laboratory.
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Si + 02 T1 + 0 2 (1) In nonpolar organic solvents, process 1 occurs with an efficiency close to unity.314 The subsequent change in the quenching oxygen molecule, on the other hand, is still uncertain. There are two important spin-allowed transitions available to the quenching oxygen molecule:s SI + 302 T1 + 302 (2)
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0022-3654/79/2083-2823$01.00/00 1979 American Chemical Society
