T H E J O U R N A L OF
PHYSICAL CHEMISTRY Registered i n U S . Patent Office 0 Copyright, 1979, by t h e American Chemical Society
VOLUME 83, NUMBER 20
OCTOBER 4,1979
Kinetics of the Reactions of the OH Radical with Hydrazine and Methylhydrazine G. W. Harris,* R. Atkinson, and J. N. Pitts, Jr. Department of Chemistry and Statewide Air Pollution Research Center, University of California, Riverside, California 9252 7 (Received February 20, 1979) Publication costs assisted by the University of California
Absolute rate constants have been determined for the reactions of the OH radical with hydrazine and methylhydrazine by using a flash photolysis-resonance fluorescence technique. The rate constants kl obtained were essentially independent of temperature over the range 298-424 K, with values of (6.1 X 1.0) X lo-'' and (6.5 f 1.3) X lo-'' cm3molecule-' s-l for hydrazine and methylhydrazine, respectively. Estimates of the lifetimes of these compounds in the atmosphere with respect to hydroxyl radical attack are made.
Introduction It has become recognized in recent years that the hydroxyl radical plays a dominant role in the chemistry of the tropo~pherel-~ and is, in many cases, the primary reactive intermediate depleting compounds introduced into the troposphere from either natural or anthropogenic sources. Hydrazine and its derivatives methylhydrazine and 1,l-dimethylhydrazine are presently used as rocket fuels, and hence the rate constants for their reaction with OH radicals are needed for any assessment of the atmospheric chemistry of these compounds following possible release into the biosphere. In this work absolute rate constants have been determined over the temperature range 298-424 K, using a flash photolysis-resonance fluorescence technique, for the reaction of OH radicals with hydrazine and methylhydrazine. Experimental Section
The apparatus and techniques used have been described previously.6 OH radicals were produced by the pulsed vacuum ultraviolet photolysis of HzOat wavelengths longer than the CaFzcutoff (11250 A). OH radical concentrations were monitored as a function of time after the flash by resonance fluorescence with a cooled EM1 9659QA photomultiplier fitted with an interference filter transmitting the 3064-A band of OH(A2C+,u ' = 0 X2H,u " = 0). The
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0022-3654/79/2083-2557$0 1.OO/O
intersection of the detection system aperture and the resonance radiation beam defined a fluorescence viewing zone at the center of the reaction vessel whose cross section was -2 cm in diameter. This region was well separated from the reaction vessel walls, minimizing heterogeneous losses of the OH radicals. The reaction cell was enclosed in a furnace which could be held constant to better than fl K over the temperature range 295-475 K, the gas temperature being measured by a chromel/alumel thermocouple mounted inside the reaction vessel. The flash lamp was typically operated at discharge energies of 10-20 J per flash at repetition rates of one flash every 3 s. Signals were obtained by photon counting in conjunction with multichannel scaling. Decay curves of OH radicals were accumulated from 118 to 2120 flashes depending on the signal strengths. OH radical half-lives ranged from 1.24 to 50.6 ms and the concentrations were followed over at least three half-lives. In all cases the flash duration (51ps) was negligible in comparison with the OH radical half-lives encountered. In order to avoid the accumulation of photolysis or reaction products and to avoid wall losses of the hydrazines (see later), we performed all experiments under flow conditions so that the gas mixture in the reaction vessel was replenished every few flashes. The partial pressure of HzO in the reaction cell ranged from 0.01 to 0.03 torr. The argon and reactants used had purity levels, according 0 1979 American Chemical Society
2558
G. W. Harris, R. Atkinson, and J. N. Pitts
The Journal of Physical Chemistry, Vol. 83, No. 20, 1979
TABLE I : Rate Constants for the Reactions of OH Radicals with Hydrazine and Methylhydrazine temp, K
reactant hydrazine methylhydrazine
298k 1 355* 1 424 f 1 298-424
500
6.5 5.9 5.8 6.5
f
t: t:
1.0 0.9 0.9 1.3
-
-
3
0
a,
v)
The quoted errors include twice the standard deviation of the slopes in Figures 1 and 3 plus the estimated errors in the measurements of flow rates, total pressures, and ultraviolet absorption calibrations,
to the manufacturers, of Ar, 199.998%; N2H4, 297% (Matheson Coleman and Bell); and CH3NHNH2,298% (Aldrich), A known fraction of the total argon flow was saturated with hydrazine or methylhydrazine vapor at 273 K, and the reactant partial pressures in this fraction of the argon flow were determined by ultraviolet absorption spectroscopy with a Cary 15 spectrophotometer. The system was calibrated by using known pressures of N2H4 and CH3NHNH2as measured by an MKS Baratron capacitance manometer. All gas flows were monitored by calibrated flow meters and the gases were premixed before entering the reaction vessel.
2 2.
a 0
w D
100
;:; 0
c
IO0
I
0
I
I
2
0
[N2H4]
I
1
4
(moiecuie
6x IO"
~ r n - ~ )
Figure 1. Plots of the OH radical decay rates against hydrazine concentration at 298, 355, and 424 K.
2 x IO+
10-11e(230i360)lRT cm3 molecule-I s-l
L
300
500
[OHl,/[OHl, = So/& = exp[(ko + kl[reactant])(t - to)] (I)
X
1
cr zoo
2 x 1
where [OH], and [OH], are the concentrations of OH at times toand t, respectively, Soand St are the corresponding resonance fluorescence intensities, ko is the first-order rate for removal of OH in the absence of added reactant (primarily attributed to diffusion out of the viewing zone and to reaction with impurities), and kl is the rate constant for the reaction of OH radicals with hydrazine or methylhydrazine. In all experiments exponential decays of the resonance fluorescence signals were observed, and the measured OH radical decay rates, defined as R = (t - to)-lIn So/& were found to depend linearly on the concentration of added hydrazine or methylhydrazine. For both reactants a variation of a factor of 2 in the flash energy produced no variation in the rate constants within the experimental errors, indicating that secondary reactions were negligible under these conditions. Furthermore, a variation from 25 to 50 torr total pressure of argon at 298 K for hydrazine and at 424 K for methylhydrazine had no effect on the rate constants within the experimental errors. Figure 1shows plots of the OH radical decay rate against hydrazine concentration for the three temperatures studied (298 f 1,355 f 1, and 424 f 1 K), while Table I gives the rate constants kl obtained from such plots by least-squares analysis. Figure 2 shows the rate constants kl for hydrazine plotted in Arrhenius form, and a least-squares analysis of the data yields the Arrhenius expression
533 400
Results The reactions of OH radicals with N2H4 and CH3NHNH2 were studied over the temperature range 298-424 K with 25-50 torr of argon as the diluent gas. Under the experimental conditions used, with the reactant concentrations being in large excess over the initial OH radical concentrations (the latter being estimated to be 110l1 molecule ~ m - ~the ) , pseudo-first-order decays of the OH radical concentrations are given by the integrated rate expression
k1(N2H4)= 4.4
1-
10"h, cm3 molecule-' s-Iu
0
-
1
0
l
I
i
1
22
,
1
26
30
,
34
I
IOOO/T(K)
Figure 2. Arrhenius plot of log of the bimolecuhr rate constant, k,(N,H,), against 1000/ T(K).
where the indicated error in the activation energy is the estimated overall error limit. However, within the error limits, the data can equally well be represented by the temperature-independent rate constant cm3 molecule-l s-l kl(N2H4)= (6.1 f 1.0) X (2' = 298-424 K) This zero or slightly negative temperature dependence (equivalent to Td.35i0.5) is in accord with the previous measurements from this laboratory on the reaction of hydroxyl radicals with a series of amines7s8 and with CH3SH7and CH3SCH2 and probably reflects a slightly temperature-dependent preexponential f a c t ~ r . ~ - ~ The only previously reported measurement of the rate constant for the reaction of OH radicals with hydrazine [kl(N2H4)= 2.2 X cm3 molecule-l s-l at 298 K'O] was obtained in a discharge flow-EPR study and is lower by a factor of -3 than the present value. The possibility that the present value of kl is high due to the participation of hydrazine-water complexes can be discounted on the basis of the study of analogous NH3.H20 complexes by Hamilton and Naleway.'l The amount of H 2 0 in the present system is very low (0.01-0.03 torr) and the ratio of its concentration to that of hydrazine varied by over an order of magnitude with no observable effect on the rate constant. Hence, the reason for the discrepancy between the present
~
OH Radical Reaction with N2H4 and CH,N2HB
The Journal of Physical Chemistry, Vol. 83, No. 20, 1979 2559
? 500
2 400
o
0
2
4
[ CH,NH, NH,]
/
6
8
i,
10 10"
(molecule crn',)
Figure 3. Plot of the OH radical decay rates against methylhydrazine at (0) 298, (+) 355,and (0) 424 K.
value and that of Hack et allo is not clear, but it is possible that heterogeneous loss processes occurring in their discharge flow system could lead to errors in understanding the stoichiometry of the reaction system. For methylhydrazine it was found necessary to allow several hours for the methylhydrazine partial pressure, as monitored in the ultraviolet absorption cell, to stabilize before commencing data acquisition. The decomposition products of methylhydrazine on surfaces are likely to include amines and, since these react rapidly with the OH r a d i ~ a l ,wider ~ ? ~ error limits were placed on the measured rate constants for methylhydrazine than on those for hydrazine. For hydrazine the decomposition products are NH3 (observed by ultraviolet absorption in the present work at the lowest hydrazine flow rates) and N2,and NH3 is less reactive than hydrazine toward the OH radical by a factor of -400.12 In view of these accuracy limits the temperature dependence of the reaction of OH radicals with methylhydrazine was investigated by measurement of only a few first-order decay rates a t the higher temperatures in order to confirm that the temperature dependence, as expected from the hydrazine data, was essentially zero. The OH radical decay rates as a function of methylhydrazine concentration are shown in Figure 3 and it can be seen that the data at the three temperatures studied are well represented by the temperature-independent rate constant k l (CH3NHNH2) = (6.5 f 1.3) X cm3 molecule-l s-l (5" = 298-424 K)
The magnitude and lack of temperature and pressure effects for these two rate constants indicate that the reactions proceed via H atom abstraction from the weak N-H bonds (N-H bond strength is -76 kcal mol-l in hydrazine13). In going from hydrazine to methylhydrazine the reaction rate with OH radicals would be expected to drop somewhat because of the fewer number of N-H bonds, but, in compensation, the electron-donating effect of the methyl group will result in a weaker bond to one of the hydrogen atoms and the net magnitude of these effects can be seen from the present work to be approximately zero. An attempt to study further these trends by measuring the rate constant for the reaction of hydroxyl radicals with 1,l-dimethylhydrazine failed because of very severe handling problems encountered for this compound, which appeared to decompose in our system very rapidly. From the rate constant data reported here for hydrazine and methylhydrazine, it may be estimated that over the temperature range -300-425 K kl(l,l-dimethylhydrazine) (5 f 2) X cm3molecule-l s-l, Using an average OH radical concentration of -lo6 cm3 for the lower troposphere, we calculated the tropospheric half-lives of all three hydrazines due to reaction with the hydroxyl radical to be -3 h.
-
Acknowledgment. The authors gratefully acknowledge the financial support of United States Air Force Contract NO. AF-F08635-78-C-0307.
References and Notes H. Levy, 11, Planet. Space Sci., 21, 575 (1973). P. J. Crutzen, Tellus, 28, 47 (1974). H. Levy, 11, Adv. Photochem., 9,369 (1974). S. C. Wofsy, Annu. Rev. Earth Planet. Sci., 4,441 (1976). P. J. Crutzen and J. Fishman, Geophys. Res. Lett., 4,321 (1977). R. Atklnson, D.A. Hansen, and J. N. Pitts, Jr., J. Chem. Phys., 82,
3284 (1975);83, 1703 (1975). R. Atkinson, R. A. Perry, and J. N. Pitts, Jr., J. Chem. Phys., 88,
1578 (1977). R. Atkinson, R. A. Perry, and J. N. Pltts, Jr., J. Chem. Phys., 88,
1850 (1978). R. Atkinson, R. A. Perry, and J. N. Pitts, Jr., Chem. Phys. Lett., 54,
14 (1978). W. Hack, K. Hoyermann, and H. Gg.Wagner, Ber. Bunsenges. Phys. Chem., 78,386 (1974). E. J. Hamitton and C. A. Naleway, J. Phys. Chem., 80,2037 (1976). R. A. Perry, R. Atkinson, and J. N. Pitts, Jr., J. Chem. Phys., 84,
3237 (1976). J. G.Calvert and J. N. Pitts, Jr., "Photochemistry", Wiley, New York,
1966.
