J. Phys. Chem. 1995, 99, 984-989
984
Atmospheric Chemistry of FNO and FN02: Reactions of FNO with 0 3 , O(3P), HO2, and HCl and the Reaction of FNO2 with 0 3 Timothy J. Wallington,* William F. Schneider, Joseph J. Szente, and M. Matti Maricq Ford Research Laboratory, Ford Motor Company, Mail Drop 3083/SRL, Dearborn, Michigan 48121 -2053
Ole John Nielsen" and Jens Sehested Section for Chemical Reactivity, Environmental Science and Technology Department, Ris@National Laboratory, DK-4000 Roskilde, Denmark Received: September 23, 1994; In Final Form: October 26, 1994@
Upper limits for the rate constants of the following gas phase reactions have been determined at 296 K: k(FNO+03) < 3 x 10-ls, k(FNO+O(3P)) < 5 x lo-',, k(FNO+H02) < 3 x 10-l2, k(FNO+HCl) < 5 x cm3 molecule-' s-'. These reactions can be ruled out as chain or sink 10-l8, and k(FN02+03) < 8 x reactions of FNO in the atmosphere, The implications for the atmospheric chemistry of FNO are discussed.
Introduction FNO is formed by the reactions of F02,' FC(0)0,233 and CF3O radicals with NO4 and is an atmospheric degradation product of hydrofluorocarbons (HFCs) and hydrochlorofluorocarbons (HCFCs).
+ NO - FNO + 0, FC(0)O + NO - FNO + CO,
(1)
CF,O
(3)
FO,
+ NO - FNO + COF,
(2)
The atmospheric fate of FNO has been the subject of recent debate.5-7 Initially, it was argued by Dibble and Francisco5 that photolysis of FNO was slow and consequently that FNO was a reservoir species for stratospheric fluorine. Wallington et aL6 commented that FNO displays significant UV absorption at wavelengths below 220 nm and in the region 280-350 nm8-12 and that photolysis is rapid. A photolytic lifetime of 12 min was estimated at 20 km,6 leading to the conclusion that the atmospheric fate of FNO is photolysis to give F atoms and NO. In response, Dibble and Francisco7 acknowledged that photolysis of FNO was indeed rapid but noted that several atmospheric species could conceivably react with FNO at rates competitive with photolysis. The reactions suggested were FNO
+ H,O -HF + HONO
AH = - 10 kcal mol-' (4)
+ O(,P) - F + NO, AH = - 17 kcal mol-' ( 5 ) FNO + 0, - FNO, + 0, AH = -22 kcal mol-' (6a) FNO + 0, - FO, + NO, AH = -4 kcal mol-' (6b) FNO + HO, - HF + NO + 0, AH = -31 kcal mol-' FNO
(7)
+
+
FNO HC1- HF ClNO AH = - 15 kcal mol-' (8) From the viewpoint of establishing the environmental impact of the atmospheric degradation of HFCs and HCFCs, reactions @
Abstract published in Advance ACS Abstracts, December 15, 1994.
4,5, and 7 produce the same result as FNO photolysis; namely, they convert FNO into HF and NO, (or HONO). The HF is inert and is transported to the troposphere where it is rained out. Reactions 5 and 6b do not impact ozone, because the lost odd oxygen is rapidly replaced by photolysis of the NO2 product. The FO2 radical produced by reaction 6b has been shown not to have any direct impact on stratospheric ozone.' Reactions 6a and 8 are interesting. Reaction 6a in concert with reaction 9 could potentially impact ozone levels via a catalytic cycle.
FNO,
+ 0, - FNO + 20,
AH = -24
kcal mol-'
(9)
Reaction 8 represents a link between FO, and C10, chemistry. To provide insight into the atmospheric chemistry of FNOx, an experimental study of reactions 5-9 has been performed in our laboratories. Reactions 6, 8, and 9 were found to proceed very slowly, if at all, and are of no atmospheric importance. Reaction 7 can at best play a minor role, whereas reaction 5 can only become important in the upper stratosphere.
Experimental Section Three experimental techniques were used to study the reactions reported herein. The long path length FTIR system at Ford Motor corn pan^'^ was employed to determine the upper limits for the FNO and FN02 reactions with ozone. Upper limits for the reactions of FNO with 0 3 , 0, and HO2 were obtained using the time-resolved flash photolysis-UV spectroscopy apparatus at Ford.14 The reaction of FNO with HC1 was examined using the pulse radiolysis instrument at F W National Laboratory.l5 FTIR-Smog Chamber System. The FTIR system was interfaced to a 140 L Pyrex reaction chamber. FNO was prepared by mixing NO with an excess of F2. FNO was then swept into the reaction chamber, which contained 0-32 mTorr of 0 3 in N2/0, diluent. Reactant loss and product formation were monitored by FTIR spectroscopy,using an analyzing path length of 28 m and a resolution of 0.25 cm-'. Infrared spectra were derived from 32 coadded spectra. FNO, FNO2, and 0 3 were monitored using their characteristic features over the wavenumber ranges 760-770 and 1830-1850, 810-830 and 1300- 1320, and 975- 1075 cm-', respectively. Reagents were obtained from the following sources and at the following
0022-365419512099-0984$09.00/0 0 1995 American Chemical Society
Atmospheric Chemistry of FNO and FNO2
J. Phys. Chem., Vol. 99, No. 3, 1995 985
purities: Fz, Matheson, >99%; NO, Matheson, >99%; NZ and 0 2 , Michigan Air Gas, >99.999%. O3was prepared by flowing 0 2 through a conventional silent electrical discharge. FNO and FNO2 were prepared by mixing NO and NO2 with an excess of Fz . Flash Photolysis-UV Spectroscopy. The apparatus consists of a reaction cell (3.2 cm diameter x 51 cm long) through which a suitable gas mixture slowly flows, an excimer laser to provide the photolysis pulse (at 193 or 351 nm), and a UV spectrometer (with diode array or photomultiplier tube detection) to record time-resolved UV spectra of the reacting gas mixture. FNO was continuously generated by flowing the desired amount of NO (0.5-2 Torr) in an excess of F2 (4-8 Torr). 0 atoms, generated in a F~/NO/OZ gas mixture by the 193 nm photolysis of 0 2 (600 Torr), subsequently add molecular oxygen to produce ozone. HOz was generated in a FflO/H2/02/N2 gas mixture by the 351 nm photolysis of FZfollowed by the reactions
+ H, - H + HF H + 0, + M - HO, + M F
(10)
(11)
Photomultiplier traces of absorbance versus time were converted to concentration profiles via Beer’s law. Time sequences of diode array spectra were deconvoluted by fitting the absorbances to reference spectra of 0 3 , HOz, and FO2, as appropriate, yielding concentration versus time profiles of the relevant species. The sources and purities of reagent gases are as for the FTIR-smog chamber system described above. Pulse Radiolysis System. FNO was produced in a 1 L stainless steel reaction cell containing 1 mbar of NO and 999 mbar of SF6 (1013 mbar = 760 Torr) using a 30 ns pulse of 2 MeV electrons from a Febetron 705B field emission accelerator:
-
+ 2 MeV e- F + products F + NO + M -FNO + M
SF,
(12) (13)
The gas mixture was analyzed using UV light from a pulsed xenon lamp. The light beam from the xenon lamp was reflected 11 times in the reaction cell by internal White type optics giving a total optical path length of 120 cm. The analyzing light was passed into a McPherson 1 m grating monochromator operated at a spectral resolution of 0.8 nm and was detected by a Hamamatsu photomultiplier coupled to a Biomation digitizer. FNO was monitored using a FNO absorption peak at 3 10.5 nm. Consistent with the published spectrum of FN0,l2 at 317 nm the absorbance was approximately a factor of 4 less than at 3 10.5 nm, indicating that FNO was indeed formed in the system. The decay of FNO in the cell was monitored with and without added HC1. The temperature was 296 & 2 K. The yield of F atoms, and thereby FNO,was determined using radiolysis of mixtures of CH4, 0 2 , and SF6 by observing the absorption at 260 nm ascribed to CH3O2 radicals formed by reactions 12, 14 and 15.3J6 In the present study the F atom yield at full radiolysis dose and 1000 mbar of sF6 was (2.8 f 0.3) x 1015 molecules ~ m - ~ .
F CH,
+ CH, -,HF + CH,
+ 0, 4- M-
CH3024- M
(14) (15)
Reagents and concentrations used were as follows: SF,5 (>99.9%), 999 mbar; CKq (>99%), 0- 10 mbar; 0 2 (99.999%), 40 mbar; HCl(>99.995%), 0-0.13 mbar; and NO (>99.8%), 1 mbar. SF6 and C& were supplied by Gerling and Holz. 0 2 was supplied by L’Air Liquide. NO was obtained from Messer
n v)
.-
+d
t 3
Time (minutes)
Figure 1. Observed decay of FNO in the presence of 750 Torr of NZ (O), 730 TOITof Nz and 20 TOITof 0 2 (A),and 730 TOITof Nz,20 Torr of 0 2 , and 32 mTorr of 0 3 (+).
Griessheim, and HC1 was supplied by Linde Techniche Gase. All reagents were used as received.
Results Reactions of FNO and FNOz with 0 3 (FTIR-Smog Chamber). FNO was prepared by addition of an excess of F2 (0.5-1 Torr) to NO (0.2-0.5 Torr) in a 2 L Pyrex bulb. Immediately following addition of F2, the mixture was swept rapidly into the reaction chamber in a stream of Nz. FNO and FNO2 were readily identified in the chamber by virtue of their characteristic IR features at 766 and 1844 cm-’ l7 and at 822, 1310, and 1796 cm-l,l* respectively. As a test for reaction of FZwith 0 3 , experiments were performed in which F2 was added to mixtures of O3in air. There was no observable effect of FZ on the O3 decay. Control experiments were performed where FNO was added to the chamber in the absence of 0 3 . Following introduction into the chamber, FNO was observed to decay with first-order kinetics as shown in Figure 1. The lifetime of FNO was approximately 10 min. As FNO decayed, FNO2 increased in concentration until all the FNO was consumed, at which time FNOz itself decayed slowly. The lifetime of FNO2 in the chamber was approximately 30 min. The reactivity of FNO toward Pyrex surfaces has been noted previously by other w ~ r k e r s . ~We ~ J ~ascribe the loss of FNO and FN02 in the chamber in NZ diluent to heterogeneous loss on the walls. When 0 2 was present in the chamber, the observed loss of FNO increased. Figure 1 shows the loss of FNO as a function of time in the presence of 20 Torr of 0 2 at 750 Torr total pressure with Nz. The first-order loss rate was 0.12 & 0.02 min-’ (errors are two standard deviations). The presence of 0 2 may increase the heterogeneous loss processes, or the walls may serve as a catalyst for the reaction of FNO with 0 2 to give FNOz. The loss of FN02 was also enhanced in the presence of 02.
Figure 2A shows the IR spectrum obtained after adding FNO to a mixture containing 20 Torr of 0 2 and 32 mTorr of O3 in 750 Torr of N2 diluent. Features attributed to FNO,FNO2, NzO5, and O3 are indicated; no other species were observed. Figure 2B shows a spectrum taken after the reaction mixture was allowed to stand for 2 min. As seen from Figure 2, FNO and 0 3 coexist in the chamber, indicating that FNO and 0 3 do not react rapidly. Figure 1 shows the observed loss of FNO as a function of time for the experiment shown in Figure 2. The FNO loss rate was enhanced in the presence of 0 3 . The enhancement may be due either to reaction 6 or to heterogeneous
Wallington et al.
986 J. Phys. Chem., Vol. 99, No. 3, 1995 I
FNO + 0,
I
I
0.8
Laser 193 nm 400 m
0.6
NO 2.0 torr F, 1.4
A:
0, 604 0.2 W
0
Z
l0.0
m
E 0
m
m
0.8
1
4
700
750
800
850
900
1000
950
1050
1100
'
WAVEN U M B E R ( c m - )
J
Figure 2. IR spectra immediately after mixing FNO sample with 03 mixture (A) and after standing in the dark for 2 min (B).
250
200
processes. The loss rate of 0 3 was unaffected by the presence of FNO and was (3-10) x min-' in the present experiments. It should be noted that the 0 3 concentration was substantially greater (by at least a factor of 12) than that of FNO. Linear least-squares analysis of the data in Figure 1 gives pseudo-first-order loss rates for FNO in the presence of 20 Torr of 0 2 diluent and 750 Torr total pressure of N2 of 0.38 =k 0.10 and 0.12 0.02 min-' with and without 31 mTorr of 03, respectively. The maximum pseudo-first-order loss attributable to reaction 6 is then 0.38 min-'. Hence, an upper limit of k6 < 7 x cm3 molecule-' s-' can be derived. In a similar fashion, an upper limit of k9 < 8 x cm3 molecule-' s-l was established from the F N 0 2 decay rate. In the above analyses it is assumed that FNO is not re-formed in any process. It might be argued that the upper limit for k6 is erroneously low because of some unknown mechanism by which FNO is generated in the system. To address this point, we can consider the observed 0 3 decay. For the experiment shown in Figures 1 and 2 the initial FNO concentration was 2.6 mTorr while the initial 0 3 concentration was 32 mTorr. After standing in the dark for 4 min, the FNO and 0 3 decayed to 28% and 96% of their initial values. An estimate of the upper limit of kg, derived from the 0 3 loss rate (1.7 x s-l) and average FNo concentration (1.7 mTorr), is k6 < 3 x cm3 molecule-' s-l. This value is entirely consistent with the value derived from the FNO decay and shows that regeneration of FNO is unimportant. When a mixture containing 1.3 mTorr of FNO2 and 30 mTorr of 0 3 (with no FNO present) was left to stand in the dark for 11.5 min, the FNO2 decayed to 59% of its initial concentration while O3 decayed to 98% of its initial value. From the 0 3 loss cm3 molecule-' we can derive an upper limit of k9 < 9 x s-l, again consistent with the value derived by considering the FNO, loss. For our final values of ks and k9 we choose to quote the lower of the upper limits obtained above, Le., ks 3 x and kg < 8 x cm3 molecule-' s-1. Reaction of FNO with 0 3 (Flash Photolysis). Although the UV cross section of 0 2 at 193 nm is very small, the photolysis of 600 Torr of 0 2 by a 400 mJ pulse produces approximately 3 x 1014 cmd3 oxygen atoms. These add 0 2
300
wavelength (nm)
0
0
100
200
300
400
500
time (ms) Figure 3. Ozone concentration versus time following the flash photolysis of 0 2 in a mixture with Fz and NO. The curve marked w/o FNO is fit to a model accounting for transport of the 03 from the photolyzed volume and its removal from the cell. The fit to the w/FNO curve includes chemical removal of 0 3 by FNO. The inset shows the absorption spectrum of the reacting gas mixture 20 ,us after photolysis.
cm6 molecule-2 S - ~ , ~ Othus with a rate constant of 6 x yielding an equal transient ozone concentration within 4 ps. Figure 3 (main figure) shows that the ozone decays both with and without the presence of FNO but that it does so more rapidly when FNO is present. Ozone is monitored instead of FNO because the latter species is in excess and has an absorption coefficient about 20 times weaker than that of 0 3 . It is the comparison of pairs of such curves, obtained under identical conditions except for the presence or absence of NO, that yields the upper limit for the FNO 0 3 rate constant discussed below. The decay in the absence of FNO is primarily due to transport of the ozone out of the volume exposed to the photolysis laser and to the slow removal of gases from the reaction cell. It is well described by a two-parameter model, one for radial transport and the other for the axial flux of material through the cell. Typical values, such as obtained for the fit to the curve marked w/o FNO in Figure 3, were 4.5 s-l and 5.5 x l O I 3 cmP3 s-l for the radial transport and axial flux, respectively; however, these varied with the gas flows used. Alternatively, the decay can be simply fit to an exponential decay with almost equal success. The addition of NO to the Oz/F2 gas mixture causes the rate of ozone decay to increase. Nitric oxide reacts rapidly with F2 ( k = 1.2 x cm3 molecule-' s-l 21) to form FNO and F atoms:
+
NO -I-F, -.+ FNO -t F
(16)
The fluorine atoms immediately add molecular oxygen to form FO2, which is then converted to FNO via reaction 1. These reactions are complete within a fraction of a second, orders of magnitude less time than required for the gases to reach the reaction cell. If one assumes, then, that reaction between FNO and ozone leads to the observed increase in the 0 3 decay rate
Atmospheric Chemistry of FNO and FNO2 and incorporates this reaction into the above transport model (or simpler exponential decay), a fit to the data yields a rate constant of (6.5 & 1.2) x cm3 molecule-' s-'. This value represents an average of five measurements with [NO] varied from 0.5 to 2 Torr and [F2]/[NO] from 5 to 10. In each case the transport and flux parameters were held fixed at the values deduced from the ozone decay in the absence of NO. Because of the very slow rate of ozone decay, it is imperative to ascertain that it is indeed due to the reaction with FNO.As mentioned above, the reactions converting NO to FNO are far too fast to allow any nitric oxide to survive transport into the reaction cell. The level of NO2 impurity in the NO sample is too small to account for the ozone decay. However, 193 nm radiation can dissociate F2 and FNO, yielding F atoms and NO. Though both react with ozone, their fate in the present gas mixture is to form FO2 and FNO,respectively, because of very high [02]/[03] and [F2]/[03] ratios. Indeed, FO2 is observed by its UV absorption, as demonstrated in the inset of Figure 3. FO2 exists as an equilibrium between F and 0 2 ; thus, the dissociation of FO2 provides a slow trickle of F atoms which are rapidly removed by reaction with ozone ( k = 1.3 x lo-" cm3molecule-' 20). The FO2 dissociationrate is of the right order to explain the increase in ozone decay caused by adding NO to the gas mixture. Unfortunately, the dissociation rate is not sufficiently well known to ascertain its contribution to the observed ozone decay; thus, the above measurement represents an upper limit summarized as k6 8 x cm3 molecule-' s-'. This result confirms the conclusion of the FTIR-smog chamber study that the reaction between FNO and 0 3 is very slow. Reaction of FNO with O(3P). A comparison of the initial ozone concentration following the photolysis of a F2/02 versus F2/N0/02 gas mixture provides an estimate for the upper limit of the rate constant for the reaction of FNO with O(3P), If oxygen atoms react sufficiently rapidly with FNO, the initial 0 3 concentration should decrease when NO is added to the F2/ 0 2 gas mixture. As Figure 3 shows, this decrease is at best a minor one ('10%). A kinetic model allowing for the competitive reaction of 0 atoms with FNO versus 0 2 shows that the upper limit for the O(3P) FNO reaction is ks 5 x cm3 molecule-' s-'. Reaction of FNO with HOz. The FNO for this reaction was continuously generated by the reaction between F2 and NO which occurs during the flow of these gases to the reaction cell. HO2 was formed by the flash photolysis of F2 by 351 nm radiation from an excimer laser followed by reactions 10 and 11. The reaction mixture was monitored by diode array UV spectroscopy. Preliminary results reveal a rapid decay of H02 and a concomitant formation of NO, as shown in Figure 4. The self-reaction of H02 is not sufficiently fast to account for its disappearance. If the reaction between FNO and HO2 is included in the model, the fits to the data illustrated in Figure 4 are obtained with a rate constant of k7 = 2 x lo-'* cm3 molecule-' s-'. A number of possible interferences preclude our assigning a definite value to the FNO HOZrate constant at this time. For example, the laser flash can photolyze FNO to produce NO, and the H atoms can react with FNO besides 0 2 . A more complete study of this reaction, which addresses these and other issues, is currently in progress. However, the observed HOz loss rate does allow us to establish an upper limit of k7 3 x cm3molecule-' s-l. (This includes the upper error bound for the previously mentioned fit.) Reaction of FNO with HCI. To determine the rate constant for reaction of FNO with HCl, mixtures of 1 mbar of NO,
+
+
J. Phys. Chem., Vol. 99, No. 3, 1995 987
5
3'0/ 2.5
0
100
200
300
400
time (ps) Figure 4. Concentration profiles of HO2 and NO following the 351 nm photolysis of a FflO/H2/0flz gas mixture.
0-0.13 mbar of HC1, and 999 mbar of SF6 were subjected to pulse radiolysis. FNO was formed in reactions 12 and 13.
FNO was monitored using its absorption at 310.5 nm. Four experiments were performed, two without HCl and two with HC1. The observed decay of FNO in the two control experiments in the absence of HC1 is shown in Figure SA. The solid lines in Figure 5A are fits of fist-order decay kinetics to the data. As seen from Figure 5A, the decay of FNO was reasonably well fit by fist-order kinetics with loss rates of (1.02 f 0.12) x s-' (filled circles) and (1.43 f 0.20) x s-' (hollow circles). As mentioned previously, FNO is reactive toward surfaces, and decay of FNO in the cell is not unexpected. Figure 5B shows the absorption at 310.5 nm as a function of time in the presence of HC1. The fist-order decay rates of FNO in these experiments were (1.41 f 0.17) x ([HCl] = 0.11 mbar) and (1.47 f 0.23) x s-' ([HCl] = 0.13 mbar). The decay rates of FNO with and without added HC1 were comparable. There is no evidence that FNO reacts with HC1. We choose to report an upper limit of k8 < 5 x IO-'* cm3 molecule-' s-' calculated by taking the average decay rate of 1.44 x s-l and [HCl] = 0.12 mbar.
Discussion The results presented herein serve to further clarify the stratospheric chemistry of FOXradicals and FNO. The known reactions of atmospheric importance for these species are summarized in Scheme 1. Photolysis of FNO produces F atoms which react rapidly with hydrogen-containing species such as C& and H20 to produce HF or with 0 2 to generate FOz radicals. HF is unreactive in the stratosphere and is ultimately transported to and rained out of the lower atmosphere. FO2 radicals, however, may re-form FNO by reaction with NO. The relative efficiencies of 02, C&, and H20 as scavengers for F atoms
Wallington et al.
988 J. Phys. Chem., Vol. 99, No. 3, 1995 50
0.04
0 (1.43io.20)x10-zs-' 0
0.03 0.02
(1.02io.12)x10-zs-1
!1 i
1
0.01
i
0.00
0
50
100
150 200 Time, s
250
300
i 0
0
2 0.05 0'06
1
w
a
v (1.47+0.23)~10~~s-'
Time, s
Figure 5. Absorption at 310.8 nm (attributed to FNO)as a function of time. Panel A shows results from control experiments in the absence of HC1. Panel B shows results when 0.11 mbar (V)or 0.13 mbar (V) of HCl was present. Solid lines are fits of first-order decays to the
data.
SCHEME 1
depend upon the altitude.' Below 4 km and above 42 km the dominant fate of F atoms is reaction with CHq and H20 to give HF.' Between 4 and 42 km reaction with 0 2 dominates. At 15 km the formation rate of FO2 radicals from F atoms is approximately 2 orders of magnitude greater than that for formation of HF, so that one FNO molecule may be photolyzed and re-formed several times before the F atom is lost as HF. A dynamic equilibrium is thus established between FNO, FO2, and F atoms, with the first two containing most of the active fluorine at steady state. The relative importance of the FNO reactions studied in this paper can therefore be determined from the ratio of the FNO removal rate by reaction with 0 3 , O(3P), HO2, or HCl (abbreviated Y) to the removal rate of F atoms by H20 and CHq (abbreviated X), i.e., from R p r o d l R ~ ~A. steady state analysis of the reactions indicated in Scheme 1 reveals that this ratio is given by
2
4
6
8
1
0
1
2
E
Figure 6. Enhancement ratio "E' as a function of altitude. See text for details. E is the number of times that FNO once photolyzed is re-formed in the atmosphere plus one.
order rate constants for reaction of F atoms with 0 2 , of F atoms with hydrogen-containing species, of FO2 radicals with NO, of FNO with 03, O(3P), H02, or HC1, and of FO2 radicals to regenerate F atoms, respectively. [FNO],, and [Flss represent the steady state concentrations of FNO and fluorine atoms, and is the FNO photolysis rate. Thus, the effect of regenerating FNO is to increase by the factor given in brackets (abbreviated E ) the branching ratio of FNO reactions with atmospheric species compared to its photolysis. Another way to think of the factor E is unity plus the number of times that FNO is reformed in the atmosphere following photolysis. The rates required to determine E have been considered previously as a function of altitude.' Inserting these rates into the above expression yields the altitude dependence of the enhancement factor, E, shown in Figure 6. Scheme 1 excludes potential minor loss processes for FO2, such as reaction with HO2 to yield HF, and assumes no decrease in kp02+N0 with decreasing temperature, both of which may serve to decrease E. More importantly, E is very sensitive to kFo2. A 10% decrease of the activation energy for FO2 dissociation from the value of 6700 K2* (used here) to the value recommended by the NASA data panel (6100 KZo) causes E to decrease by roughly a factor of 10. Thus, the enhancement of the FNO branching ratio shown in Figure 6 is expected to be an upper limit. It shows that regeneration of FNO is r.iaximized at an altitude of about 15 km and that at this altitude FNO is expected to be regenerated fewer than 10 times. In the stratosphere the maximum concentrations of 0 3 , HO2, and HC1 are approximately 6 x 10l2, 2 x lo7, and 3 x lo9 ~ m - respectively, ~, whereas the O(3P) concentration drops from 7 x lo9 cm-3 at 50 km to about 6 x lo5 cmP3 at 15 km.20 Upper limits of k5 < 5 x k6 < 3 x k7 < 3 x and k8 < 5 x cm3 molecule-' s-' at 296 K are reported herein. Typical stratospheric temperatures are much less than 296 K, and it is reasonable to assume that the room temperature upper limits obtained here will be valid at the lower stratospheric temperature also. The minimum lifetimes of FNO with respect to reactions 6-8 are then 5.6 x lo4, 1.7 x lo4, and 6.7 x lo7 s, respectively; for O(3P) it increases from 2.9 x lo2 s at 50 km to 3.3 x lo6 s at 15 km. The lifetime of FNO with respect to photolysis in the stratosphere is approximately 720 s . ~ In the worst case, the photolysis lifetime is effectively increased relative to the FNO reaction lifetime by the enhancement factor E given in Figure 6, Le., by about a factor of 10 at 15 km and decreasing to unity at 50 km. Even with the
Atmospheric Chemistry of FNO and FN02
J. Phys. Chem., Vol. 99,No. 3, 1995 989
conservative estimates employed here, we find that reactions 6 and 8 cannot compete effectively with FNO photolysis in the stratosphere. Using the pseudo-first-order rate constants appropriate to 15 km described above, we obtain RFNO+O,IRHF < 0.13. In the worst possible case at this altitude reaction with ozone is 11% as effective in removing FWO as is photolysis and formation of HF. Similarly, RFNO+HC~/RHF < 1 x and reaction with HCl can be excluded as a significant loss process for FNO. The FNO HO2 reaction could be marginally competitive, but only at the 15 km peak of the FNO regeneration. At this altitude the upper limit for the probability of reaction with HO:! is RFNO+HO-JRHF .e 0.43. Even if competitive, however, this reaction yields the same products as FNO photolysis followed by the reaction of F with a hydrogen-containing species; thus, its impact on stratospheric chemistry will be negligible. With the upper limit of k5 obtained here, the FNO O(3P) reaction could become competitive in the upper stratosphere. At 50 km, and ignoring temperature effects in reaction 5 , we obtain RmO+dRm = 2.5, so that reaction with O(3P) atoms may dominate the chemistry of FNO at this altitude. However, the NO2 product would be rapidly photolyzed to replenish the odd oxygen species removed by this reaction. The remaining reaction suggested by Dibble and Francisco’ as potentially important in the atmospheric chemistry of FNO is reaction with H20. N o data are available for this reaction. Both FNO and H:!O are closed-shell species, and to postulate a rapid reaction between these two seems unreasonable. As noted above, reaction 4 converts FNO to HF and thus has the same effect as photolysis of FNO. Finally, the potential for heterogeneous chemistry of FNO needs to be considered. For typical ambient stratospheric conditions the collision frequency of gas molecules with aerosol surfaces is approximately s-1,23 and typically only 1 collision in 10 leads to reaction?O In the stratosphere, photolysis of FNO proceeds approximately 1 order of magnitude faster than collisions with particles. Hence, heterogeneous processes will be of minor importance. In conclusion,it is reasonable to consider photolysis the major
+
+
atmospheric fate of FNO,although further work is necessary to clarify the roles of reaction with HOz radicals and O(3P) atoms. Such work is in progress in our laboratories.
Acknowledgment. We thank Ted Dibble (Wayne State University) for helpful discussions. References and Notes (1) Sehested, J.; Sehested, K.; Nielson, 0.J.; Walhgton, T.J. J . Phys. Chem. 1994, 98, 6831. (2) Wallington, T.J.; Ellermann, T.; Nielsen, 0. J.; Sehested, J. J . Phys. Chem. 1994,98,2346. (3) Maricq, M. M.; Szente, J. J.; Dibble, T. S.; Francisco, J. S. J . Phys. Chem. 1994, 98, 12294. (4) Chen, J.; Zhu, T.; Niki, H. J . Phys. Chem. 1992, 96, 6115. (5) Dibble, T. S.; Francisco, J. S. J. Phys. Chem. 1994, 98, 5010. (6) Wallington, T.J.; Schneider, W. F.; Nielsen, 0. J.; Sehested, J. J. Phys. Chem. 1994, 98, 10373. (7) Dibble, T. S.; Francisco, J. S. J. Phys. Chem. 1994, 98, 10374. (8) Johnston, H. S.; Bertin, H. J. J . Mol. Spectrosc. 1959, 3, 683. (9) Solgadi, D.; Flament, J. P. Phorophys. Photochem. 6 eV 1985,497. (10) Suter, H. U.; Huber, J. R.; von Dirke, M.; Untch, A,; Schinke, R. J. Chem. Phys. 1992, 96, 6727. (11) Ogai, A.; Brandon, J.; Reisler, H.; Suter, H. U.; Huber, J. R.; von Dirke, M.; Schinke, R. J . Chem. Phys. 1992, 96, 6643. (12) Burley, J. D.; Miller, C. E.; Johnston, H. S. J . Mol. Spectrosc. 1993, 158, 317. (13) Wallington, T.J.; Japar, S. M. J . Armos. Chem. 1989, 9, 399. (14) Maricq, M. M.; Wallington, T. J. J . Phys. Chem. 1992, 96, 986. (15) Nielsen, 0. J. Rise-R-480, 1984. Ellermann, T. Rise National Laboratory Report Rise-M-2932, 1991. (16) Sehested, J. Int. J . Chem. Kinet. 1994, 26, 1023. (17) Woltz, P. J. H.; Jones, E. A.; Nielsen, A. H. J . Chem. Phys. 1952, 20, 318. (18) Bernitt, D. L.; Miller, R. H.; Hisatsune, I. C. Spectrochim. Acra 1967, 23A, 237. (19) Flores, A. L.; Darwent, B. beB. J . Phys. Chem. 1969, 73, 2203. (20) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R.; Kolb, C. E.; Molina, M. J. Jet Propulsion Laboratory Publication 92-20, Pasadena, CA, 1992. (21) Westley, F.; Herron, J. T.;Cvetanovic, R. J.; Hampson, R. F.; Mallard, W. G: NIST Chemical Kinetics Database, NIST, Gaithersburg, MD, 1991. (22) Hippler, H. Private communication, 1993. (23) Turco, R. Rev. Geophys. Space Sci. 1982, 20, 233.
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