1931
REACTION OF CYANOGEN RADICALS WITH AMMONIA less likely t o occur than some of the others. The same is true to a lesser extent for reaction 15. Reactions 13 and 14 appear most favorable for SZOformation.
Conclusion Obviously, a number of reactions exist which could be important in the formation of the intermediates SZ, SzO, and SO. Many of thc rcactions involve radicals and atoms essential to flame propagation. The three sulfur-bearing intermediates from H2S oxidation are, therefore, likely to form generally in any flame oxidation process. From the air pollution point of view, the removal of any one of the three intermediates in a combustion process would, of course, reduce SO2 formation in the system. However, since the formation of SO2 is believed t o occur predominantly through the reaction
so +
0 2
=
so2 + 0
(16)
it is apparent that removing so Jvould greatly reduce SO2formation and, therefore, its emission to the atmosphere. There are obvious problems involved in any attempt to remove any of the intermediates-problems such as finding reactants which will react rapidly and specifically with S2, S20, or SO and which will form compounds with the intermediates that will withstand the high temperatures encountered in flame processes. Kevertheless, the possibility of removing the sulfurbearing intermediates to control sulfur oxide emission should not be ignored.
Acknowledgment, The authors wish to acknowledge the assistance of the Research Grants Branch, Environmental Protection Agency, under Grants AP 00464-04 and -05.
The Reaction of Cyanogen Radicals with Ammonia1 by G. E. Bullock, R. Cooper,* Chem&ry Department, University of Melbourne, Parkville, Victoria 5052, Australia
S. Gordon, and W. A. Mulac Chemistry Division, Argonne National Laboratory, Argonne, Illinois
60489
(ReceiPed J a n u a r y 10, 197.8)
Publication costs assisted by Argonne National Laboratory
The pulse radiolysis of dilute C2N2-NH3 mixtures in argon has been used t o study the reaction of CN radicals with NH, at 300 and 375’K. An increase in the rate constant for CN radical disappearance was observed for radicals with higher vibrational excitation than the ground state. Removal of the ground and fourth vibrational states showed a negative temperature dependence. (‘(Theseresults are discussed with respect to the possible roles played by chemical reaction and by vibrational relaxation.”)
Introduction Cyanogen (CN) radicals have received a great deal of attention in spectroscopic2 and thermodynamic3 studies but fewer data are available on the kinetics of reactions of this reactive radical. Bimolecular rate constants of varying reliability have been determined for reaction with oxygenj4J ammonia,5 methane,5 nitric and the dimer cyanogen (C2N2).4a~G Some measurements have included estimates of activation s5,6
The results of work done at this laboratory on the rates of reaction with saturated and unsaturated hydrocarbons are to appear shortly? I n this paper we wish to report kinetic parameters for the reaction of CN radicals with gaseous ammonia a t 303 and 375°K. CN
radicals, produced by pulse radiolysis of dilute solutions of CzWzin argon, have been monitored spectrophoto(1) Based on work performed under the auspices of the U. S. Atomic Energy Commission. (2) (a) G. Herzberg, “Molecular Spectra and Molecular Structure, Vol. I. Spectra of Diatomic iMolecules,” 2nd ed, Van Xostrand, Princeton, N. J.. 1950. and references therein: (b) N. Basco. J. E. Nicholas, R. G. W. Norrish, and W. H. J. Vickers, Proc. R o y . Soc., Ser. A , 272, 147 (1963). (3) (a) D. D . Davies and H. Okabe, J . Chem. Phys., 49, 5526 (1968); (b) T. L. Cottrell, “The Strengths of Chemical Bonds,” Butterworths, London, 1954. (4) (a) D. E. Paul and F. W. Dalby, J. Chem. Phys., 37, 592 (1962); (b) N. Basco and R. G. W. Norrish, Proc. R o y . Soc., Ser. A , 283,291 (1965); Ii. Basco, ibid., 283, 302 (1965). (5) J. C. Boden and B. A. Thrush, ibid., 305, 107 (1968). (6) C. A. Goy, D. H. Shaw, and H. 0. Pritchard, J. P h y s . Chem., 69, 1504 (1965).
T h e Journal of Physical Chemistry, V o l . Y 6 , N o . 14, 19YR
1932
G. E. BULLOCK, R. COOPER,S. GORDON, AND W. A. MULAC
metrically in the ground, second, and fourth vibrational levels of the ground electronic state (XzZ). An attempt was also made to follow the appearance of the product NHZ radical.8
I
I
I
I
I
I
I
I
1
I
Experimental Section The pulse radiolysis equipment has been described in detail previously.8 The radiolysis cell used in this work, while of the same design, was modified in several ways t o allow controlled elevated temperatures to be obtained. “Viton” O-rings were replaced by gold and the adjustable mirror controls were changed t o a temperature-insensitive type using a metal bellows rather than O-rings. A thermocouple (Pt-10% Rh-90% Pt) was included to monitor the gas temperature. The cell was completely enclosed in a thermostated oven which permitted accurate temperature control ( f~ 1 ~ 5 ° K up to -425°K. The optics and kinetic spectroscopic arrangements were unchanged from the earlier publications.7r8 The Hilger-Engis Model 600/1000 vias used as a monochromator and with 100-p entry and exit slits gave a band pass of -0.8 b. This meant the vibrational bands of the CN X22--+ B28electronic transition could be separately monitored via the Av = 0 sequence at -3880 A. These wavelengths were arbitrarily chosen but were -1.5 8 below the band heads. The absorptions did not obey Beer’s law and the modified form, OD = ~ ’ ( d ) ” was , u ~ e d . 7 The ~ ~ n values were determined by varying either path length or dose7 and were effectively constant at 0.75 f 0.05 for the three bands. For the NH, radical a n n value of 0.77 was used.8 The cyanogen was Matheson Co. CP grade and was purified as described previ~uslp.~ Ammonia was stored over metallic sodium and purified by conventional trap-to-trap distillations. Argon was Matheson Ultra-Pure and was subject to freeze-thaw-pump cycles in the vacuum system prior t o use. The cell mas typically filled with 15 Torr of C D Z , 0-3 Torr of NH3,and 700 Torr of argon. The ammonia pressure was either measured directly or via an ammonia-argon mixture which gave higher and therefore more accurately measurable pressures. The differences here were insignificant.
Results The removal of CK radicals in the absence of scavengers other than the parent C2Nz occurred relatively slowly. The observed first-order rate constants with 15 Torr of C2X2 and 700 Torr of argon were < 2 X lo4 sec-l and 50 times more efficient than CH, as a vibrational energy scavenger in these systems. Further, thc relaxation of the fourth level of CN by SH3 must be occurring on approximately every twentieth collision, an efficiency beyond the upper limit suggested by Lambert14 for a nonresonant process. Callcar and Williarnsls observed fast relaxation of S O by a number of triatomic hydrides and suggested11’15that the anomalous efficiency may be due to hydrogen bonding in the collision complex (see later). (ii)Influence of Vibrational h’neryy on Reaction Rates. Considerations of vibrational energy in chemical kinetics have in the past concentrated on two cases: the accumulation of energy in unimolecular reactions and the production of vibrationally excited species as a result of reaction. Despite the absence of any preccdcnt,l6 it remains possible that internal energy in the form of vibration of one of the reactants may be able to contributc to the overall energetics of reactions. Transition state theory predicts that variations in vibrational partition functions on forming a transition state can cause a rate enhancement. Unfortunately, such predictions are doubtful when the activation energy is negligibly small and -it appears that this is the case in this reaction (see later). The NH2 appearance was followed in an attempt to determine the overall reaction rate constant. As shown (Figure 4 and Table I), the rate constant, -1.15 X 10’0 1. mol-’ sec-l, is in good agreement with the value for CK(0,O) removal. Apart from the larger errors in these data, the high intercept in Figure 4 causes some doubts about the validity of the rate constant obtained from the slope. It implies that there
+
T h e Journal of Physical Chemistry! Vol. 7 6 , S o . 14, 1.973
is a rapid production of NH2 which does not come from reaction 3 .
Temperature Effects For the high-temperature work (Figure 3 ) , the concentrations of NH3 were calculated at the temperature measured by the thermocouple in the gas itself. Any effect of nonequilibration should thus be eliminated. It seems that the lower rate constant at higher temperatures is real and we conclude that there is a small negative activation energy of
