4159
COMMUNICATIONS TO THE EDITOR
Among the species which may contribute to the spectra of Figure 1 are X, Xz-, X2+,RXaX, R X f , and RX-. The esr spectrum of gaseous I atoms1° extends from 3300 to 6223 gauss. We find a spectrum extending from 1500 to 5000 gauss from yirradiated polycrystals prepared by fusing KI with 0.1 mole % AgN03, a medium in which the optical spectrum has been attributed to IZ-.l1 The esr spectrum of Br2- in Xirradiated KBr single crystals containing alkaline earth ions is observed12between 2000 and 4800 gauss. Samples were irradiated to a dose of 4 X l O l 9 ev g-’ a t a dose rate of 2 X 10ls ev g-’ min-l. Magnetic field values were assigned with a Varian Fieldial accessory. The modulation amplitude of 4000 (about 12 gauss) used gave a maximum esr signal without distortion. The other experimental methods used have been described.“
modes of forming or breaking bonds accessible, in principle, to such compounds as N20, 0 2 , and GO2. Nevertheless, certain trial postulates (outlined below) have been formulated which provide remarkably good agreement between computed and measured Eo, and, accordingly, show promise of adding to our understanding of the nature of bimolecular reactions. These Eo computations do not use adjustable parameters, but rely on bond properties such as dissociation energy and vibrational frequency. For the multivalent bimolecular B =A BX, one of the trial transfer reaction AX postulates relating the bond order nz (in the transition state A. * .X. .B) of X . . B to the bond order nl of A . - + Xcan be written
+
n2
+
= (n”Bx/n”Ax)(n”Ax
- nl)
(1)
where n”Ax is the bond order of the reactant AX, and n ” B X is the bond order of the product BX. Equation 1 (10) S. Aditya and J. E. Willard, J . Chem. Phya., 44, 833 (1966). meets the requirement that nl is n”Ax when n2is zero (11) C. J. Delbecq, W.Hayes, and P. H. Yuster, Phys. Rev.,1 2 1 , at the initiation of the reaction, and vice versa. For 1043 (1961). univalent reactions, eq 1 reduces to the previously (12) W. Hayes and G. M. Nichols, ibid., 117, 993 (1960). used1~2~4 trial postulate that nl n2 = 1. CorrespondDEPARTMENT OF CHEMISTRY RICHARD J. EGLAND 1 is equivalent to the simple postulate that ingly, eq UNIVERSITY OF WISCONSIN JOHN E . WILLARD the bond-order increase of the forming bond is linearly MADISON, WISCONSIN 53706 related to the decrease in bond order of the breaking RECEIVED JULY17, 1967 bond, with the proportionality constant equal to the ratio of the bond order of BX to that of AX. As in univalent calculations, the potential energy of activaV , in the transition state is taken as the energy tion, Computed Activation Energies for Bimolecular lost by the bond AX as it dissociates to A . . . X of bond Reactions of 0 2 , Nz, NO, NzO, NOZ,and C02 order nl, less the energy supplied by the formation of B . . . X of bond order n2, plus a repulsive energy V , Sir: Although it had once been hoped that absolute arising from parallel electron spins on -4and B. rate theory would permit useful kinetic predictions for = D e , ~x D’Ax(ni)lrAXmost reactions, the problem of predicting activation energies is still one of the major unsolved questions in D ’ x B ( ~ ~ ) *V~, kcal/mole ~ (2) chemistry since tractable accurate quantum-mechanical where D e , ~ Xis the dissociation potential energy of AX, solutions do not exist for potential energy surfaces of D’ is the dissociation potential energy of the single reactions involving multielectron atoms. For hydrobond, and V , is obtained from eq 10 of ref 4. The gen atom reactions in the gas phase, it has been shown, A is log (D,/D’)/log n”, which is the slope exponent however, that transition-state computations’, of activaof the log dissociation energy us. log bond order line. tion energies and rate constants by a nonquantum n1 or n2 in eq 2 is less than one, A is replaced by When method can exhibit good agreement with experiment the slope p , calculated’ for the bond-order region below when several tentative assumptions are used to provide one. As in previous investigation^,'^^^^ the computer interrelationships (such as Pauling’s rule3) among bond program (modified for multivalent bonds) determines energy, bond order, and bond length. In an effort to avoid this computing method’s limitation to H-atom (1) H. S. Johnston and C. Parr, J . Am. Chem. Soc., 85, 2544 (1963). reactions, a reduced-variable treatment was devised4 (2) S. W. Mayer, L. Schieler, and H. S. Johnston, J . Chem. Phys., for other univalent atoms, such as the halogens. 45, 385 (1966). The calculation of activation energies for the im(3) L. Pauling, “The Nature of the Chemical Bond,” 3rd ed, Cornel1 University Press, Ithaca, N. Y . , 1960. portant class of bimolecular reactions involving multi(4) S. W. Mayer, L. Schieler, and H. 5. Johnston, “Proceedings of valent bonds has been an even more difficult challenge the Eleventh International Symposium on Combustion,” The Combecause of the complexities introduced by the many bustion Institute, Pittsburgh, Pa., 1967, p 837.
+
v
+
Volume 71, Number 1.9 November 1967
COMMUNICATIONS TO THE EDITOR
4160
the value of nz corresponding to the peak in the potential energj', V , of the transition state. The computed Eo is equal to V plus a small zero-point energy correction. Activation energies computed on the basis of eq 1 and 2 and data in the JXSAF thermochemical tables are summarized in Table I and 11, where the reactions are written in the exothermic direction. Although the
Table I : Comparison of Computed with Measured Activation Energies
Reaction
AX
+ I3 = A + BX
Rfeasured
Eo
Computed Eo, kcal/moIe Spin Ground conserwstate tion spin
2 0 0 37 0 1 24
50a 27a 15a 505 0s 32b 27" 7c
...
05
...
48 29
15 49
1 25 29
8 0
Reviewed by K. Schofield, Planet. Space Sci., 15, 643 (1967). R. M. Fristrom and A. A. Westenberg, "Flame Structure," McGrawHill Book Co., Inc., New York, N. Y., 1965, p 371.
* H. 5.Johnston, et al., J . Chem. Phys., 26, 1002 (1957).
Table 11: Transfers between Two Free Radicals, Negligible rlctivation Energya
.kX
ONO
Reactionb = A
+R
+ Ii
+ BX
= ON
+ OH
+ 0 = O N + 02 ON + N = 0 + Nz I10 + 0 := H + 02 O N + 0 == 0 + N O + 0 = 0 + 0, ON0
0 2
Transition state formula
HONO ON02 N20
HOz ON0 0 3
Computed activation energy, kcal/ mole
EnergyC of new bond, kcalj mole
1 3 0 5 1 4
49 (61) 63 72 26
78
(See J. T. Herron and F. S. a The experimental Eo is -0. Klein, J. Chtm. Phys., 40, 2731 (1964) and source cited in footnote a of Table I.) b Spin conservation in the transition state can occur here via the ground state species. c JANAF thermochemical tables.
thermal dissociation of XZO probably involves singlet O(lD) crossin@ to trip1et 0(3p) primarily, the 'Omputed Eo's of Table 1 indicate that bimolecular transiThe Journal OJ' Physical Chemistry
tion states for NzO reactions are consistent with W i p e r spin conservation since the trial postulate of spin conservation dissociation of NzO(lZ) in the transition state through O(lD) and Nz('Zg+) leads to very good agreement between computed and measured Eo, whereas the calculation of Eo for the reaction mechanism allowing crossing to ground state O(3P) produces very poor agreement with experiment. Furthermore, this spin conservation (and noncrossing) postulate for bimolecular transition states apparently also extends to the reactions involving COZ(lZ) and OH(Q), since Eo computed for the transfer of O(lD) generally agrees much better with experiment than does Eo computed for crossing to O(3P). All of the reacbions in Tables I and I1 exhibit conservation of orbital angular momentum6 although momentum conservation is generally less restrictive than spin conservation. On the basis of the results summarized in Table 11, another trial postulate can be suggested. Consider, for example, the transfer reaction between NO2 and atomic hydrogen. Both of these reactants have an unpaired electron. When the H atom approaches the S O z radical, bond formation between the H atom and an oxygen atom on XOZ can begin, therefore, without requiring corresponding breaking of the bond between W and 0. If the structural formula of the transition state is similar to that of a stable molecule (ie., in which BX would have De> 25 kcal/mole), the energy released during this stage of forming the new bond can provide the activation potential energy required by eq 2. The experimental Eo could be negligible in such cases, since tjhe energy released during the early stage of bond formation between the free radicals would be subtracted from the computed Eo. The results in Table I1 support the postulate that experimental Eo will be negligible for exothermic transfers in which the transition-state structure meets the aforementioned conditions. For each of these six reactions, the observed Eo is less than 1.0 kcal/mole, and De for BX is sufficiently high so that the computed requirement for Eo could readily be supplied by the initial bond formation between the reactant free radicals. No exceptions to this postulate have been found. The spin repulsion term, V,, of eq 2 was generally small, less than 2 kcal/mole, for the reactions of Tables I and 11. A fuller examination of these bimolecular transitionstate postulates of spin conservation and stable bond formation will be prepared, along with an evaluation (5) E. K. Gill and K. J. Laidler, Can. J. C'hem., 36, 1570 (1958). (6) K, J , Laidler, ,'The Chemical Kinetics of Excited States," Oxford University Press, Oxford, England, 1955, p 22.
COMMUNICATIONS TO THE EDITOR
4161
of the justification of the assumptions inherent in eq 1 and eq 2. Acknowledgment. It is a pleasure to acknowledge the valuable correspondence with Professor H. S. Johnston of the University of California, Berkeley, Calif. LABORATORIES DIVISION AEROSPACECORPORATION EL SEGUNDO, CaLIFoRNIA 90245
S. W. MAYER
RECEIVED AUGUST25, 1967
t i L
!
235
I
I
26 0
255
I - & 310
WAVELENGTH IN MILLIMICRONS
Radiolytic Products of Liquid Ammonia'
Sir: Radiolysis of liquid ammonia results in the formation of Hz, N2, and N2H4.'-' A recent publication reports that HN3 is also a radiolytic product.' In this communication, we present the results of experiments8 showing that a t least two other products, one of which is "3, are formed in the radiolysis of liquid ammonia. Samples of liquid ammonia that had been purified by storage over sodium, followed by trap-to-trap distillation, were irradiated at 10" with Co60 y-rays for total doses of 1 X lo4 to 1.5 X 108 rads. The yields of H2, Nz, and N2H4 were measured. The amount of excess hydrogen, calculated from the material balance equation H2,excess
=
Hz,obsd - NZH4,obsd - 3N2,obsd
was plotted as a function of dose. Two distinct dose regions, (1) I 3 0 X lo3 rads, where G(H2,excess) = 0.51 f 0.03, and (2) >50 X lo3 rads, where = 0.14 0.03, were found where there G(H2,excess) was a lack of material balance. Thus either additional products with a N2:HZ ratio > 1:2 were present or impurities were an important influence in the radiolysis. No evidence for any impurity was found. When alkaline aqueous solutions of ammonia, which had been irradiated to doses 5 3 0 X lo3 rads, were examined spectroscopically, two broad absorption bands centered at 285 and 243 mp were observed (Figure 1, curve 111). On acidifying, these bands disappeared and a more intense band at 230 mp was observed (Figure 1, curve IV). These absorption spectra were similar to that of a control solution of tetramethyl tetrazene (curves I and 11) and to that previously observed for alkyl tetra~enes.9~'~These results, together with the observation that on increasing the dose these bands disappear with the formation of nitrogen, suggest that the unknown species contains a K=N group. Possible compounds which may be present are , (N3H3), or tetrazene (X4H4). diazene (SzHz)triazene
Figure 1. Typical ultraviolet spectra of aqueous solution of liquid ammonia irradiated to doses 5 3 0 X 103 rads: I, solution of tetramethyl tetrazene, p H 2.8; 11, solution of tetramethyl tetrazene, p H 11.8; 111, aqueous solution of irradiated ammonia, p H 11.8; and IV, aqueous solution of irradiated ammonia, p H 2.8.
Typical ultraviolet absorption spectra of aqueous solutions of ammonia, which had been irradiated to doses >50 X lo3 rads, are shown in Figure 2. In basic and neutral solution (curve 111), absorption spectra show little structure, while in acid solution (curve IV), there is a broad absorption band a t 260 mp. This behavior is similar to that of control samples of NaN3 (Figure 2, curves I and 11). The presence of N3- was confirmed by measuring the infrared spectra of (1) an acidified aqueous solution of irradiated NHs and (2) of a KBr pellet prepared from t.his solution (Table I). Absorption bands characteristic of azide were observed. A further experiment was carried out in which helium was bubbled through an acidified aqueous solution of irradiated liquid ammonia and passed through a liquid nitrogen trap. The trap was allowed to warm up and the resulting vapor expanded (1) This work was performed under the auspices of the U. S. Atomic Energy Commission. (2) D. Cleaver, E. Collinson, and F. S. Dainton, Trans. Faraday SOC.,56, 1640 (1960). (3) L. Kolditz and U. Prosch, Z . Physik. Chem., 208, 108 (1962). (4) J. R. Puig and E. Sehwars. "Industrial Uses of Large Radiation Sources," Vol. I, International Atomic Energy Agency, Vienna, 1963. (5) F. S. Dainton, T. Skarski, D. Smithies, and E. Wemamowski, Trans. Faraday SOC.,60, 1068 (1964). (6) D. Schischkoff and D. Schulte-Frohlinde, 2. Phgsik. Chem., 44, 112 (1965). (7) J. Belloni, J . Chim. Phys., 9, 1281 (1966). (8) J. W. Sutherland and H. Kramer, Annual Reports, Nuclear Engineering Department, Brookhaven National Laboratory, Upton, Long Island, N. Y.: BNL 900 (5-67). p 83, 1964; BNL 954 (5-681, p 88, 1965; BNL 994 (AS-20), p 61, 1966. (9) T. M.Bins and N. R. McBride, Anal. Chem., 31, 1382 (1959). (10) N. R. McBride and €1. W. Kruse, J . Am. Chem. SOC.,79, 572 (1957).
Volume 71,Number 19 November 1967
