3308
J. Phys. Chem. 1984, 88, 3308-3313
at the high EA values of A and C. This is in contrast to the BSBL method, which would not predict major differences in reactivity among the three C1 reactions or the three O H reactions, and the formulae of Alfassi and Benson, which would predict near-zero activation energies for all six reactions. The latter occurs because the majority of the reations used by these authors" to establish their empirical fit involved combined EA'S of less than 3 EV. Those reactions in the 3-4.5-eV range of combined EA had widely scattered activation energies (0-4 kcal/mol). Thus we would not expect this empiricism to be adequate for the reactions discussed here. We can, however, make a number of qualitative observations based on this data. First, differences in C1(2P) atom reactivity (EA = 3.61 eV) far exceed the differences in OH(X211) reactivity (EA = 1.83 eV) for the set of three reactions involving reactants substituted with varying strength electrophiles. For both sets, the reaction rate decreases with increasing electron affinity of the substituent group. In the case of OH, the changes in reeactivity, as mirrored in the room temperature rate constants, correlate well on a plot of log k vs. EA of the substituent. The similarities in activation energies for the O H CH3Cl and OH CH3CN reactions possibly reflect the near equal electron affinities of the C1 and C N substituents. The lower rate constant for OH CH,CN arises as a result of a lower Arhenius A factor which we noted earlier as indicative of a tight activated complex or one of restricted geometry. In this light one might speculate that the extremely large dipole moment in CH3CN (see Table 11) causes a rather tight alignment in the HO-CH,CN reaction complex (pOH = 1.7 D), resulting in a more negative AS of activation. In the C1 atom reaction series, the large decrease in reactivity from CH3C1to CH3CN is disproportionate relative to the small change in substituent EA. Again, it is possible that a strong electrostatic interaction between the electrophilic C1 and the CH3CN dipole leads to a very restricted reaction-complex geometry which is not conducive to H atom transfer. While this in part would be manifested by a somewhat higher activation energy than observed for CH3Cl, one might also expect such reductions in transitionstate entropy to result in a considerably lower Arrhenius A factor. Such interactions would be expected to have similar qualitative effects for reactions proceeding via a complex pathway involving a bound intermediate. Tight geometries for such an adduct could preclude rapid rearrangement necessary for product formation. (This interpretation was dismissed for reaction 1 due to the sizable positive activation energy observed.) Although the above observations and speculations are far from being quantitative, they do point out certain inadequacies in our present abilities to predict chemical reaction, parameters. The
+
six reactions we have selected for advancing such speculation were chosen due to similarities (within each set of three) in exothermicities and bond dissociation energies. Under these constraints the body of gas-phase literature data available for comparative purposes is quite limited. Clearly, far more experimental data are needed to test the validity of the hypotheses proposed here and to develop more quantitative correlations. For example, it would be of interest to see how the effects of such electronic interactions are partitioned between the activation energy and preexponential factor, how they are affected by the characteristics of the atoms or group transferred, how they are influenced by BDE's or reaction thermochemistry, etc. As a final point of discussion we can consider the effect of the present results on atmospheric models. First, because of the lower rate constant and higher activation energy measured in this work for reaction 1, calculations of OH mixing ratios in the 33-45-km region (based on the observed CH3CN profile and the assumption of dominant CH3CN loss by reaction 1) are in better agreement with OH measurements.2 This moderate change in rate constant does nothing, however, to alleviate the major discontinuity between ground level concentrations3of CH3CN and the lower stratospheric datae2 A major tropospheric loss for CH$N (possibly washout) must still be hyp~thesized.~ The absence of significant reaction between CH3CN and atomic chlorine dispels any suggestion that reaction 2 could play a dominant role as a loss of acetonitrile or chlorine atoms in the upper stratosphere. While we might reasonably conclude that loss coefficients for homogeneous atmospheric reaction as well as photodissociation are adequately specified, considerable uncertainity remains in elucidating the severe altitudinal variation in CH3CN. Brasseur et al.,5 for example, calculate that the ground level flux of CH3CN required to account for these measurements far exceeds the sum of natural and anthropogenic NO, sources. The minor modeling changes caused by the revisions in reaction rates from the present work do not alter these conclusions. Thus it is of critical importance to understand the nature of any heterogeneous removal of CH3CN during tropospheric transport.
+
+
Acknowledgment. This work was supported in part by the Fluorocarbon Program of the Chemical Manufacturers Association and the Upper Atmospheric Research Program of the National Aeronautics and Space Administration. The authors thank Drs. L. W. Sieck and R. F. Hampson of NBS for helpful discussions during the preparation of this manuscript. Registry No. CH,CN, 75-05-8; atomic chlorine, 22537-15-1; hydroxyl, 3352-57-6.
Absolute Rate Constants for the Reaction of Hydroxyl Radicals with Ammonia from 297 to 364 K Robert D. Stephens Environmental Science Department, General Motors Research Laboratories, Warren, Michigan 48090 (Received: November 15, 1983)
-
Absolute rate constants for the reaction OH + NH3 NH2 + H 2 0 were measured as a function of temperature from 297 to 364 K via the discharge flow technique with resonance fluorescence detection of OH. The rate constant was measured exp[(-973 f 78)/T]/cm3 molecule-' at five temperatures and fit to an Arrhenius expression to yield (4.55 f 1.1) X s-I. A computer model of the experimental data was used to estimate rate constants for the exothermic reactions NH2 + NO N2 + H + OH and NH2 + NOz N2 + 20H. The rate constant for quenching of OH(A2Z+)at 297 K by ammonia cm3 molecule-' s-'. was determined t o be (9.6 f 0.8) X
-
-
Introduction Ammonia is the third most abundant nitrogen species in the atmosphere. It is thought to play a significant role in both homogeneous and heterogeneous atmospheric acid neutralization 0022-3654/84/2088-3308$01.50/0
reactions.',2 However, possibly as much as 50% of atmospheric ammonia is oxidized in the gas phaseS3x4The primary reaction (1) S. c. Wofsy and
M. B. McElroy, Can.J . Chem., 52,
0 1984 American Chemical Society
1582 (1974).
The Journal of Physical Chemistry, Vol. 88, No. 15, 1984 3309
Reaction of OH with Ammonia
Reactant
responsible for the initial step in ammonia oxidation is4 NH3
+ OH
NH2
-+
+ H20
1
(1)
Pressure
II
Multiport Valve
Several measurements of the rate constant for this reaction have been reported in the literature ranging from 4.1 X to 2.7 x 10-l~cm3 molecule-' S-I at room t e m p e r a t ~ r e . ~ - ' ~ The ultimate fate of the amidogen radical (NH,) thus formed is unclear, however. It has been implicated as a potentially significant source for atmospheric NO, via the reaction^'^^^^ N H 2 + O2 NH2
+ O3
--
products
(2)
products
(3)
Unambiguous rate constant measurements for reaction 2 do not exist, but upper limits have been reported in the range 8.3 X to 2 x cm3 molecule-' s-' a t room temperature and a b ~ v e . " ~ ' ~Rate J ~ constants for reaction 3 have been reported cm3 molecule-' s-l at room in the range 6.3 X to 1.2 X temperature.I5-l7 The amidogen radical could also be a significant sink for atmospheric NOx1,4,6via the reactions
--
+ NO NH2 + NO2 NH2
products
(4)
products
(5)
The rate constant for reaction 4 has been reported in the range of 1.2 X lo-" to 2.7 X lo-'' cm3 molecule-' s-l at room temperature,'8-22 whereas reaction 5 has a rate constant measured to be 1.2 X lo-" to 2.3 X lo-" cm3 molecule-' s-1.19,24 Several possible exothermic channels exist for reactions 2-5, but few product identifications have been made. It is likely that reactions 3 and 4 occur by more than one pathway. For example, OH has been detected a t levels of a few percent of NH2 in experiments performed on reaction 3,16 and perhaps 20% of reaction 4 proceeds via an OH formation process.'8 The formation of OH as a product of reaction 5 is also thermodynamically possible. Clearly, further study is required to elucidate the rates and mechanisms of potential atmospheric reactions of NH,. Given the rate constant measurements for reactions 1-5 and the approximate atmospheric concentrations of OH, NH3, and NO,, one can conclude that the rate-limiting step in ammonia (2) T. E. Graedel, "Chemical Compounds in the Atmosphere", Academic Press, New York, l'978. (3) "Report of the NASA Working Group on Tropospheric Program Planning", NASA Reference Publication 1062, 1980. (4) J. C. McConnell, J . Geophys. Res., 78, 7812 (1973). (5) I. W. M. Smith and R. Zellner, Znt. J . Chem. Kine?. Symp., I , 341 (1975). (6) F. Stuhl, J. Chem. Phys., 59, 635 (1973). (7) R. A. Perry, R. Atkinson, and J. N. Pitts, Jr., J . Chem. Phys., 64,3237 (1976). ( 8 ) S . Gordon and W. A. Mulac, Znt. J. Chem. Kine?. Symp., 1, 289 (1975). (9) J. A. Silver and C. E. Kolb, Chem. Phys. Lett., 75, 191 (1980). (10) V. W. Hack, K. Hoyermann, and H. Gg. Wagner, Ber. Bunsenges. Phys. Chem., 78, 386 (1974). (1 I ) P. B. Pagsberg, J. Eriksen, and H. C. Christensen, J. Phys. Chem., 83, 582 (1979). (12) M. J. Kurylo, Chem. Phys. Lett., 23, 467 (1973). (13) R. Lesclaux and M. Demissy, Nouv. J . Chim., 1, 443 (1977). (14) S. G. Cheskis and 0. M. Sarkisov, Chem. Phys. Lett., 62,72 (1979). (15) H. Kurasawa and R. Lesclaux, Chem. Phys. Lett., 72,437 (1980). (16) W. Hack, 0. Horie, and H. Gg. Wagner, Ber. Bunsenges. Phys. Chem., 85, 72 (1981). (17) V. P. Bulatov, A. A. Buloyan, S. G. Cheskis, M. Z. Kozliner, 0. M. Sarkisov, and A. I. Trostin, Chem. Phys. Lett., 74, 288 (1980). (18) J. A. Silver, C. M. Gozewski, and C. E. Kolb, ARI-RR-216, Western States Section 1 Combustion Institute, Oct 20, 1980. (19) W. Hack, H. Schacke, M. Schroter, and H. Gg. Wagner, (Znt.) Combust., [Proc.],17th, 505 (1978). (20) G. Hancock, W. Lange, M. Lenzi, and K. H. Welge, Chem. Phys. Lett., 33, 168 (1975). (21) R. Lesclaux, P. V. KhB, P. Dezauzier, and J. C. Soulignac, Chem. Phys. Lett., 35, 493 (1975). (22) S. Gordon, W. Mulac, and P. Nangia, J . Phys. Chem., 75, 2087 (197 1). (23) K. Schofield, J . Chem. Phys. Ref.Data, 8, 723 (1979). (24) H. Kurasawa and R. Lesclaux, Chem. Phys. Lett., 66, 602 (1979).
Wood's Horn
It'
-v
Microwave Cavity
He"20
\ Pump Resonance Lamp
Figure 1. Schematic diagram of discharge flow resonance fluorescence system.
oxidation is reaction l.I2 It is, therefore, important that the rate constant for this reaction be accurately established. The initial objective in studying NH, chemistry was, therefore, to unambiguously ascertain the rate constant of this reaction as a function of temperature. Experimental Section
The discharge flow apparatus is similar to that used previously25 and is illustrated in Figure 1. It consists of a resonance fluorescence detection cell and a removable 2.5-cm i.d. X 3.8-cm 0.d. Teflon-coated aluminum tube. Eight fixed-position reactant inlet tubes in. 0.d.) were equally spaced from 10 to 45 cm from the detection zone and were positioned to open at or near the center of the axis of major gas flow. Other ports were used for pressure and temperature measurements as well as for a side arm through which radicals were introduced. The flow tube and fluorescence cell were wrapped with an electrical heating element and insulation. A movable thermocouple probe (retracted during kinetics measurements) was used to verify that temperature control was maintained to within f 2 K throughout the reaction zone. Hydroxyl radicals were produced 30-cm upstream of the first reactant inlet port by reacting atomic hydrogen (produced by microwave discharge through 92-ppm hydrogen in helium) with an excess of nitrogen dioxide. Concentrations of nitrogen dioxide were chosen such that hydroxyl radical formation was always complete (Le., three H atom half-lives) within 3 ms. Hydroxyl radicals were detected downstream via resonance fluorescence induced by a resonance lamp consisting of a microwave discharge through approximately 2 torr of l% H 2 0 in helium. Fluorescence signals were collected at right angles to the fluorescence excitation source through an interference filter and detected via a photomultiplier (EM1 9789QA) operated in the photon-counting mode. Apertures and focusing optics were utilized with both optical axes to maximize the ratio of signal to scattered light. Hydroxyl radical sensitivity was estimated at 3 X lo9 molecule cm-3 at a signal-to-noise ratio of one. Concentrations of hydroxyl radicals were measured daily by titration of H atoms with nitrogen dioxide and never exceeded 6 X 10" molecules ~ m - ~ . The clean Teflon-coated surface of the flow tube was found to have unacceptably high and irreproducible wall loss rates for hydroxyl radicals. Coating the flow tube and injectors with halocarbon wax (Halocarbon Products Corp. Series 15-00) minimized this problem and was used for all experiments. Wall loss rates for hydroxyl were measured daily by varying the point of nitrogen dioxide injection and determining the slope of In (OH fluorescence) vs. wall exposure time. Wall loss rates measured in this way never exceeded 6 s-'. This technique assumes instantaneous hydroxyl formation, no H-atom wall loss, and no hydroxyl loss due to reaction with excess nitrogen dioxide. The first two assumptions lead to an underestimation of hydroxyl wall ( 2 5 ) L. G. Anderson and R. D. Stephens, 15th Informal Photochemistry Conference, Stanford, CA, 1982, p 201.
3310
The Journal of Physical Chemistry, Vol. 88, No. 15, 1984
Stephens
I
KOBS (s-') vs NH3 CONC
-2
2oo! 150
I
0
5
I
I
1
10 16 Reaction Time ( G e c )
I
20
Figure 2. Plots of the natural log of OH fluorescence signal vs. reaction time at 334 K. The curves shown correspond to 1.61 X 1014,3.01 X lOI4, and 5.14 X 1014 molecules of ammonia ~ r n - respectively, ~, from top to bottom. Error bars represent statistical uncertainties of 2a.
01
0
I
I
I
I
I
I
1
2
3
4
5
6
Ammonia, 1 O1
loss. This underestimate is partially offset by the reaction of hydroxyl with nitrogen dioxide which, under the conditions employed, occurs at approximately 1 s-l. Hydroxyl wall loss rates were also determined from the intercepts of plots of pseudofirst-order decay rates of O H fluorescence vs. reactant concentration. Hydroxyl wall loss rates determined in this manner never exceeded 13 8. Gases used were helium (Matheson U H P 99.999%), hydrogen in helium (92-ppm Matheson certified standard), nitrogen dioxide in helium (274 ppm and 0.917% Matheson certified standard), and ammonia (Matheson anhydrous 99.99% minimum). The 0.917% nitrogen dioxide in helium was analyzed via long-path Fourier transform infrared spectroscopy and found to contain 550 ppm of nitric oxide as a contaminant. All gases were used directly without further purification. Pressure in the flow system was measured at 50 cm from the detection zone by an MKS capacitance manometer. Temperature was measured with type-T thermocouples. All gas flow rates were controlled and measured with Tylan mass-flow controllers calibrated by measuring the pressure rise resulting from a timed flow into a known volume. The kinetics experiments were performed under pseudo-firstorder conditions by adding ammonia in large excess over hydroxyl, sequentially through each of the eight reactant inlet ports. Measurements of scattered light and O H fluorescence in the absence of ammonia, IFo, were made both immediately before and after the kinetics experiments were conducted. Measurements of IFowere used in the determination of the rate constant for quenching of OH(AZB+)by ammonia as discussed later. Acquisitions of fluorescence count rates, gas flow rates, pressure, and temperature were accomplished by a CAMAC standard kinetic systems 8010 microcomputer system. Statistical tests were applied to these measurements to verify that data collection occurred during stable experimental conditions. A series of eight measurements of each parameter was made at each reaction time and the data were remeasured if the standard deviation in any measurement was greater than 3%. Remeasurement of data was typically not required under these conditions. The measured pressure and calculated flow velocity were corrected for viscous pressure drop to derive the pressure and velocity at the point representing the average of the midpoints between reactant injection and hydroxyl detection. These corrections were typically approximately 5%. The corrected flow velocity (typically 20 m s-') and pressure ( 1 torr) were used to calculate the reaction time and species concentrations (based on the fraction of total flow measured for each gas). N
Results OH NH,. Reaction 1 was studied at five temperatures from 297 to 364 K at 1 torr. With [NH,] > 1.4 X 1014molecule cm-3 >> [OH] I 6 X 10" molecule ~ m - reaction ~, 1 is pseudo-first order
+
molecule
Figure 3. Plot of pseudo-first-order rate constant for OH fluorescence decay vs. ammonia concentration. Data are for 334 K.
1245L
-12 5 0
7
-1255
r
I \
-12.60
4
s
-12.65
~
m
-12.70
- 1 2.75
' 0
-12.80 I
-12 85 -12901 275
I
I
I
I
2875
30
3 125
3 25
I
1 /T ( x 1000). K - l
Figure 4. Arrhenius plot for the temperature dependence of OH + NH3 reaction rate constant. Error bars represedt statistical uncertainties of 2a.
in OH. Figure 2 shows typical plots of In [OH fluorescence] vs. reaction time. Slopes of such plots yield a value for the observed pseudo-first-order rate of OH decay, KOM,where KoM = kl [NH,] k , and k , is a first-order term for removal of OH at the flow tube wall. Intercepts of such plots yield a value for the fluorescence intensity, IF,in the presence of ammonia but in the absence of O H removal by reaction with ammonia. The values for IFwere employed in a Stern-Volmer analysis as discussed later. A plot of Kob&vs. [NH,] is shown in Figure 3. The value for kl at each temperature is derived from the slope of such plots as calculated by a weighted least-squares fit of the data. The value for k l at 297 K determined in this way is (1.73 0.11) X cm3 molecule-' s-I, A summary of Kobrd,temperature, and ammonia concentration for each experiment, as well as the values of k l , is listed in Table I. The values for k l for the five temperatures were fit to an Arrhenius expression to yield a temperature-dependent rate constant of (4.55 f 1.0) X exp[(-973 f 7 8 ) / q cm3 molecule-' s-l. The stated uncertainties in all cases represent one standard deviation. The Arrhenius plot is shown in Figure 4. OH(A*Z+)Quenching. From Figure 2 it can be noted that the O H fluorescence signal at zero reaction time varies as a function of ammonia concentration. Under ideal conditions resonance fluorescence would result from two processes:
+
*
-
+ OH OH(A22+) OH(AZB+)2OH(X2rI,) + hv' hv
(a) (b)
The Journal of Physical Chemistry, Vol. 88, No. 15, 1984 3311
Reaction of OH with Ammonia
1014
temp, K
molecules cm-3
Kobsd,(l S-1
297 297 297 297 297 297 297 320 320 320 320 320 320 320 334 334 334 334 334 334 334 334 348 348 348 347 347 347 348 363 363 363 363 363 363 363 363
2.66 4.33 6.24 7.91 9.68 11.4
57.0 f 1 88.5 f 2 130 f 2 148 f 2 168 f 4 195 4
*
k l , cm3 molecule-l
56.8 f 1 84.0 f 2 105 f 2 123 f 1 140 f 2 158 f 1
1.61 2.30 3.01 3.71 4.43 5.14 5.86
49.9 f 1 68.2 f 2 87.0 f 1 104 f 2 120 f 1 136 f 2 156 f 2
1.55 2.14 2.67 3.21 3.13 4.24
48.9 f 1 64.3 f 1 80.9 f 2 101 2 114 f 2 119f 1
1.47 1.98 2.50 3.00 3.54 4.10 4.59
32.2 f 1 46.6 f 1 59.6 f 1 86.1 f 1 98.4 f 1 111 f 1 125 f 1
-U -2 Y
* 0.06) x 10-3
Uncertainties are one standard deviation. where hv' is the OH fluorescence to be detected and k, is the rate constant for the radiative decay of OH(A2Z+). However, the nonradiative process of collisional quenching can also occur: OH(X2Hi)
+M
(c)
where kqMis the rate constant for the quenching process via molecule M. Fluorescence intensity is, then, a function of quenching. Under the experimental conditions employed, M is predominantly helium, NO2, and NH3. The OH(A2Z+) fluorescence lifetime, T ~ O ,in the absence of ammonia is
= (k,
+ kqHe[He]+ kqNoz[N02])-'
The OH(A2Z+)fluorescence lifetime in the presence of ammonia, T F , is TF
= (k,
+ kqHe[He]+ kqNo2[N02] + kqNH3[NH3])-l
At a given hydroxyl concentration, the ratio of OH fluorescence to the fluorescence intensity in the absence of ammonia, IFo, intensity in the presence of ammonia, IF,is IFO/IF=
TF'/TF
=1
I
/
0
I
I
I
20
I
I
I
I
NH3 Concentration
I
I
80
60
40
I
l
100
1
l
120
Mole~ule/cm~~)
Figure 5. Stern-Volmerplot of 297 K data showing IFO/IFvs. ammonia concentration. Error bars represent statistical uncertainties of 20.
(3.07 f 0.15) X
7F0
1.3-
1.11
(2.74 f 0.12) x 10-13
kqM
1.4-
. - I/
*
-
15-
i n
(2.48 f 0.03) X lo-"
+M
1.71.6-
(2.12
OH(A22+)
1 .E-
s-I
(1.73 i 0.11) x 10-13
2.1 1 3.34 4.18 5.05 . 5.92 6.99
I
191
TABLE I: Compilation of Experimental Data from k l Determination -INHl1, __
+ TFO~,"'[NH~]
This is the typical Stern-Volmer quenching expression which can be employed to derive a rate constant for the quenching of OH(A2Z+)by ammonia. The values of IFused for these calculations were obtained in conjunction with the kl determination by extrapolating the In (OH fluorescence) vs. reaction time data (Figure 2) to zero reaction time. This provides a measure of OH fluorescence both in the presence of ammonia and in the absence of removal of OH by reaction with ammonia. The values for IFo were the fluorescence signals measured in the absence of added
ammonia. Ammonia flows used for the k , determination never exceeded 3% of total gas flow; hence, concentrations of other quenchers remained essentially constant between measurements of IFoand experiments resulting in derivation of IF. When the radiative lifetime of OH(A2X+) and the rate constants for quenching of OH(A2Zt) summarized by Schofield are used,23 the OH(A2Z+) fluorescence lifetime in the absence of ammonia ( T ~ O ) is essentially equal to the radiative lifetime of OH(A22+), s. Given the gas flow velocities employed in i.e., -0.76 X the experiments (-20 m s-') and the spatial resolution of fluorescence detection (-0.5 cm), the time resolution of fluorescence detection is long relative to processes a-c, hence these processes are in steady state. A least-squares analysis of the data (Figure 5 ) yields a value for 7FokqNH3 of (7.3 f 0.6) x cm3 molecule-'. Hence, the 297 K rate constant for OH(AZZ+) quenching, as measured cm3 relative to the radiative lifetime, is (9.6 f 0.8) X molecule-I s-l. This is in good agreement with the value reported by Fairchild et a1.26and could be interpreted as support for their postulated temperature dependence of quenching phenomena. This rate constant is approximately four times greater than the 297 K collision rate constant for these species and is only exceeded by the reported quenching rate constants for CC14 and CFC13.28 Although possible, there is no evidence that quenching involves chemical reactions. This possibility was not investigated in these experiments. It should also be noted that the measurement of kl and k,, although subject to the same uncertainties, are each independent of the other.
Discussion Data from experiments performed at room temperature and high ammonia concentrations displayed a slight nonlinearity in plots of In (OH fluorescence) vs. reaction time as shown in Figure 6. Although the data are linear within the statistical uncertainties, the slight nonlinearities are suggestive of an effect of secondary chemistry. The potential fate of N H 2 formed via reaction 1 was therefore investigated. Recent studies have revealed three possible branches for reaction 4: NH2 N O N2 H 2 0 (4a)
+
+ NO NH, + N O
NH2
- ++ + - +
+
N2
H
N2H
OH
OH
(4b)
(4c)
One study indicates that approximately 20% or more of reaction (26) P. W. Fairchild, G. P. Smith, and D. R. Crosley, J . Chem. Phys., 79, 1795 (1983). (27) P . Andreson, A. Jacobs, C. Kleinermanns, and J. Wolfrum, Sym. (Int.) Combusr., [Proc.], 19rh, 11 (1982). (28) M. A. A. Clyne and P. M. Holt, J . Chem. SOC., Faraday Trans. 2, 75, 569 (1979).
3312 The Journal of Physical Chemistry, Vol. 88, No. 15, 1984
Stephens
1
TABLE II: Chemical Model Used To Examine the Perturbation on Measured [OH] Resulting from Reaction 4b or 5b'
26 -
C
E
-
reaction rate constant OH + NO 1.23 X cm3 molecule-I s-' NH2 + 1.7 X lo-" cm3 molecule-' s-' N2 + 0-2.2 X lo-" cm3 H2O (4a) molecule-'
+ NO2 OH + NH3 H20 NH2 + NO H
25-
0 C
c"
-
V
2
-
24-
NH2 + NO N2 + H + OH (4b) NH2 + NO2 N 2 0 + H2O (sa)
-C
23
_-
-
0
NH2 + NO2 N2 + 20H (5b) NH2 + H + M NH3 +M NH2 + NH2 + M .--* N2H4 + M NH2 + OH NH + H20 OH + NO2 + M HNO, + M OH + NO + M -, HONO + M OH + OH H2O + 0 +
5
10
15
Reaction Time
20
25
(ms)
Figure 6. Comparison of experimental data (normalized to measured OH
concentration)and the computer model predictions of OH concentration for 297 K. Solid lines represent the weighted least-squares fit of experimental (@). The computer model predictions of OH concentration (X) were obtained by using kdb= 5 X cm3molecule-' s-' and k5b = 0 cm3 molecule-' s-'. Similar model predictions result with kSb= 5 X cm3molecule-' s-' and k4b = 0 cm3 molecule-' s-'. The curves represent ammonia concentrations (from top to bottom) of 4.33 X lOI4, 7.91 X lOI4, 9.68 X lOI4, and 1.14 X 10l5molecules cnr3. 4 yields OH; however, no N2H was detected.I8 A separate study concludes that greater than 69% of reaction 4 produces OH, yet no H atoms were detected.27 In these experiments, reaction 4b would result in the production of two O H for each NH, reacted, since O H would also be formed via the fast reaction of H with the nitrogen dioxide in the system. There is evidence that reaction 5 occurs predominantly but not exclusively vial9 NH2 NO:, N 2 0 H20 (5a)
+
+
+
However, there is an exothermic channel for this reaction which could also result in OH formation: NH2 + NO2 N2 + 2 0 H (5b) The potential impact of reactions 4b and 5b on measured O H decays in this study would therefore be nearly identical, Le., two OH formed for each NH, reacted. Reactions 4 and 5 have both been measured to have negative temperature dependences which would be consistent with the observation of nonlinearity of experimental data at low temperature only. Furthermore, this nonlinearity would be expected to be most prominent under conditions of high N H 2 concentration, which is consistent with its occurrence only in experiments performed at high ammonia concentrations. To determine the possible influence of reactions 4b and 5b on the measurements reported herein, a computer model of the dominant chemistry present in the experiments was used to simulate the results at 297 K. The chemical reactions and rate constants used in the simulation are shown in Table 11. The simulations were performed with a program written by WhitbeckZ9 that uses the technique described by Gear.30 Simulations were based on experiments at room temperature and utilized the measured [OH], k,, [NO,], and pressure. To determine the potential influence of reaction 5b on the measurements reported herein, the rate constants for (5a) and (5b) were varied while maintaining the overall rate constant at k5 = 2.2 X lo-" cm3 molecule-l s-' and assuming k4b = 0. The model predicted the cm3 molexperimental Kobsdfor hydroxyl when k5b= 5 X ecule-' d. The model predicted OH concentrations outside the +
+
+
-+
-
OH + wall NH2 wall
0-2.2 x lo-'' cm3 molecule-' s-' 2.9 X cm6 molecule-2s-' 6.9 X cm6 molecule-2 s-I 5.0 X lo-" cm3 molecule-l s-' 2.6 X lO"O cm6 molecule-2s-l 6.7 X cm6 molecule-2 s-I 1.8 X cm3 molecule-' s-I 12.6 s-I 12.6 s-'
this work
sum of 4a and 4b set at average of ref 20, 21, and 22 sum of 5a and 5b
set at average of ref 19 and 24 33 33
14
32 31
32 this work assumed equal to k, for OH as measured in this work
"The initial concentrations were as follows: [NO] = 1 X 10l2molecules cm-3;b[NO,] = l X l O I 3 molecules ~ m - ~[NH,] ; = 4.33 X 10i4-1.14 X 1015 molecules ~ m - [MI ~ ; = 3.2 X 10l6 molecules cm3; [OH] = 4 X 10" molecules c ~ r - ~bBased . on the measured NO contamination of the NO2 used, plus the NO resulting from the OH formation reaction. 20 statistical uncertainties of the O H fluorescence measurements when kSb = 1 X lo-'' cm3 molecule-' s-'. To determine the potential impact of reaction 4b on the measurement of kl the rate constants for (4a) and (4b) in the simulations were also varied while maintaining the overall rate constant at k4 = 2.2 X lo-'' cm3 molecule-' s-' and assuming k5b = 0. The simulations indicate that the experimental results were relatively insensitive to the branching ratio of reaction 4 since the dominant loss of NH, was via reaction 5. However, the best match between the experimentally obtained Kobsd and Kobsdobtained from the simulation occurred when reaction 4b was predicted to proceed at approxcm3 molecule-' s-I, Le., approximately 23% of imately 5 X reaction 4. This is in reasonable agreement with existing branching ratio measurements. The simulation results are shown with experimental data in Figure 6. The rate of O H formation predicted by the model from either reaction 4b or 5b is less than 10% of the rate of OH loss due to reaction 1 throughout the range of ammonia concentrations utilized in the experiments. Furthermore, the model predicts OH concentrations within the 2g limits of statistical uncertainty of the experimental data when kdb and k j b are both zero. The modeling of these experiments therefore suggests that O H formation from N H 2 reactions is possible if not probable.35 However, the data demonstrate that k l is not significantly perturbed by this secondary chemistry. Of seven studies of kl previously reported (summarized in Table 111), four are in reasonable agreement and yield the preferred value for k , of 1.6 X cm3 molecule-' s-l at room t e m p e r a t ~ r e . ~ ~ ~~
(29) M. R. Whitbeck, private communication. (30) C. W. Gear, 'Numerical Initial Value Problems in Ordinary Differential Equations", Prentice-Hall, Englewood Cliffs, NJ, 1971.
2.2 X 10-"-0 cm3 molecule-' s-I 2.2 X 10-"-0 cm3 molecule-' s-I
ref 32
~
~
(31) H. Okabe, "Photochemistry of Small Molecules", Wiley, New York, 1978. (32) R. F. Hampson, "Chemical Kinetic and Photochemical Data Sheets for Atmospheric Reactions", D.O.T. Report No. FAA-EE-80-17, 1980.
The Journal of Physical Chemistry, Vol. 88, No. 15, 1984 3313
Reaction of OH with Ammonia
TABLE III: Comparison of Results of This Study to Previously Reported Values of k l
room temp rate constant, cm3 molecule-’ s-l
A , om3 molecule-’
1.57 X lo-”
SKI
2.3 X lo-’,
E , cal mol-’ 1600
1.5 ( f 0 . 4 ) X lo-”
228-472 298
1.64 (f0.16) X
2.93 X lo-’,
1.44 (f0.29) X lo-”
(5.41
2.42 x 10-13
(5.3 f 0.8) X lo-’,
j:
0.86) X lo-’,
1.1 x 10-12
(2.7 f 0.3) X
1710 f 300
297-427
2120 f 143
294-1075
1830
298-669
870 f 80
(4.1 f 0.6) X
(1.73
temp range, K
298-365 298
0 . 1 ~ )x 1043
(4.55 f 1.1) x 10-1,
1933 f 155
Three of four flash photolysis studieP7 and one of two discharge flow studies9 support this result. A potential for secondary chemistry in the flash photolysis studies exists due to the potential for ammonia photolysis with subsequent reaction of the N H 2 photofragment via
NH2
+ OH
-
products
(6)
The rate constant for reaction 6 has been estimated at 5 X lo-” cm3 molecule-’ s-l.14 Since the absorption cross section for ammonia, across much of the ultraviolet spectrum, is equal to or greater than that for water3’ (the OH source in most flash photolysis studies), the concentration of NH2 produced by both photolysis and reaction 1 could be expected to exceed the concentration of OH. Hence the rate of OH decay in the flash photolysis studies might be expected to increase beyond that resulting from reaction 1. However, Smith and Zellner5 carefully checked for and eliminated this possible complication in their study. The agreement obtained in two of three other flash photolysis studies which did not investigate this possibility suggests that the rate constant for reaction 6 has been overestimated. The possible complicating secondary chemistry in the discharge flow studies results from the potential for OH regeneration via reactions 4 and 5 as discussed previously. This would result in a decreased rate of OH decay from that expected’dueto reaction 1. Of the two discharge flow studies, only Hack et al.,1° considered secondary chemistry. Their result, however, does not agree with the flash photolysis studies. The discharge flow study of Silver and Kolb9 did not investigate secondary chemistry. Furthermore, the temperature range of their study is not comparable with the (33) V. A. Lozovskii, V. A. Nadtochenko, 0. M. Sarkisov, and S. G. Cheskis, Kinet. Katal., 20, 1118 (1979). (34) D. L. Bauch, R. A. Cox, P. J. Crutzen, R. F. Hampson, Jr., J. A. Kerr, J . Troe, and R. T. Watson, J. Phys. Chern. Ref.Data, 11,393 (1982). (35) In fact, preliminary experiments performed with this discharge flow apparatus, using NH2 formed via the reaction of F atoms with NH,, have confirmed the formation of OH (as detected via resonance fluorescence) from either reaction 4 or 5. Subsequent detection of nitric oxide contamination of the nitrogen dioxide used in these experiments precludes the confirmation of reaction 5 as the sole source of this OH, however. Further investigations of these reactions are currently in progress.
297-364
method flash photolysis resonance absorption flash photolysis resonance fluorescence flash photolysis resonance fluorescence discharge flow resonance fluorescence discharge flow electron spin resonance pulsed radiolysis kinetic spectroscopy flash photolysis resonance fluorescence discharge flow resonance fluorescence
ref 5 6 7 9 10 11 12
this work
flash photolysis studies, except at room temperature where their result is in agreement. This study, then, offers an alternative to the flash photolysis technique while also carefully considering potential secondary chemistry. The good agreement obtained between this study and most flash photolysis studies (despite the potential for secondary chemistry which could effect the results in opposite directions) offers strong confirmation for the preferred value of k l . As reaction 1 is the rate-limiting step in the atmospheric gas-phase oxidation of ammonia, it controls the extent to which the tropospheric NO, budget can be influenced by reactions 2-5. The extent of this influence could far exceed NO, sources due to combustion.6 An understanding of this influence awaits knowledge concerning the relative importance of reactions 4 and 5 (NO, sink) as opposed to reactions 2 and 3 (NO, source), as well as the relative importance of reaction 1 to heterogeneous or other ammonia loss processes. Further investigations of the rates and mechanisms of reactions 2-5 are, therefore, still required. The technique employed to concurrently obtain the rate constant for OH(A2Z+) quenching and OH(X*II,) reaction kinetics has not, to the author’s knowledge, been previously employed. A more cumbersome technique to accomplishe this has been reported by Clyne and Holt.26 However, it is worthwhile to note that the technique reported herein can be easily applied to much of the extensive data in existence for ground-state radical reaction kinetics, with no further measurements required. Furthermore, this technique is equally valid for flash photolysis experiments. Care must be taken, however, to ensure stable intensities in the fluorescence excitation and the photolysis source so as to maintain constant concentration and excitation of OH(X211i) between measurements of IFand IFo.This means that the reactant must not significantly absorb either wavelength employed.
Acknowledgment. The author expresses his gratitude to Dr. Larry G. Anderson for his preparation of software used for data acquisition and analysis as well as for many helpful discussions, to Dr. David Dolson for his helpful review of this manuscript, and to Dr. Jerry D. Rogers for the gas analysis. Registry No. OH, 3352-57-6; NH3, 7664-41-7; NH,, 13770-40-6; NO, 10102-43-9; NO,, 10102-44-0.
