1687
Nanosecond Pulse Radiolysis of Hydrazine
Nanosecond Pulse Radiolysis of Hydrazine J. DeIalre,* P. Cordier, J. Belloni, F. Billiau, and M. 0. Delcourt Laboratoire de Physico-Chimie des Rayonnements,Laboratoire associe au CNRS, 9 1405 Orsay, France (Received February 4, 1976) ~
Publication costs assisted by Laboratolre de Physico-Chlmie des Rayonnements, Orsay
The transient spectrum obtained in pure liquid hydrazine shows a broad absorption extending from 500 to 1100 nm which has been attributed to the solvated electron; its maximum wavelength is located beyond 1100 nm. A small uv band is also found, probably due to oxidizing species. The recombination rate constant of the solvated electron with the acidic ion N2H5+ is (5 f 3) X 107 M-l s-1 and hence is not diffusion controlled. The electron yield (G = 3.4) in the nanosecond time range has been obtained by scavenging with biphenyl. The rate constant of the scavenging reaction is 1.7 X 1010M-l s-l and the extinction coefficient of the solvated electron is e = 1.85 X lo4 M-l cm-l at 1000 nm. The high value of electron yield is interpreted in terms of the slowness of the recombination with the cation NzH5+.
Introduction Theoretical models based on diffusion attempt to predict the fate of radiation induced species as a function of time and space, and of parameters related to the initial space distribution and to the physical properties of the solvent such as diffusion coefficients, dielectric constants, and relaxation times. In fact diffusion models imply that the solvated electron only reacts with the cation and that this neutralization effectively occurs at each encounter. This assumption could be incorrect and experimental results on liquid ammonia afford an example which favors consideration of collision efficiency. Indeed the yield of the ammoniated electron14 is much higher than expected by comparing with the electron yields obtained from solvents of similar static dielectric constant. We explained2 this peculiarity by the marked slowness of the neutralization of the ammoniated electron by the acidic cation NH4+ (K = lo6 M-ls-15). This explanation is also that given by Farhataziz, Perkey, and Hentz4 and confirmed by their results on the electron yield dependence on pressure.6 The aim of this paper is to extend the “G value-rate constant” correlation to hydrazine. Its physical properties are comparable to those of water whereas its chemical properties7 resemble those of liquid ammonia (amine function and the proton affinity). A few results on the radiolysis of liquid hydrazine have been obtained by y irradiationaIO and by pulsed techniquell but no experimehts have been carried out in the nanosecond time range. Experimental Section The radiation source was a 706 Febetron delivering a 3-115 pulse of 600-keV electrons and providing about 1014electrons per pulse. The pulse radiolysis set-up will be described in detail elsewhere.12Because of the catalytic decomposition of hydrazine by metals, instead of using a stainless steel cell,l3 an entire silica cell has been designed14with a 0.2-mm silica window for entrance of the electron beam. The optical path length is 1cm and the cross section of the analyzing beam is 0.05 X 1 cm2.The fast spectroscopicand electronic equipment enables us to detect transient absorbing species from 200 to 1100 nm and to record kinetics with a 3.7-ns risetime. A light beam splitting technique is used for determination of spectra, except for biphenyl experiments for which the spectrum has been normalized with respect to that of the solvated electron.
Hydrazine was distilled over quicklime under reduced pressure of inert gas according to the procedure described in ref 15. Hydrazine contained less than 0.5% water and 0.3% ammonia which corresponds to a negligible concentration of basic and acidic ions. Hydrazinium chloride was from Prolabo. Zone refined biphenyl was from Merck. Hydrazine samples and solutions were degassed by the freezepumpthaw technique (77 K, Torr). Samples were irradiated at room temperature and repetitive pulsing had no effect on results. The experimental dose was 4.7 X 1017 eV ml-l per pulse as determined from the initial absorption of the hydrated electron, taking G,,,- = 3.316 and eeaq- = 13 000 M-I cm-l at 600 nm. Biphenyl concentrations were measured spectrophotometrically after extraction by hexane from hydrazine solutions. Results 1. Transient Spectra and Kinetic Data. Figures 1 and 2 show the transient absorption spectra obtained in pure liquid hydrazine at the end of a 3-1-15pulse and a t time 4 ps after irradiation. Transient spectrum was not extended below 310 nm because of the strong absorption of hydrazine. For the wavelength range 450-1100 nm, a similar spectrum has been previously observed by the use of a 4.5-ps electron pulsell and has been attributed to the solvated electron in hydrazine. Evidence for such an assignment results from scavenging experiments in the presence of biphenyl and hydrazinium chloride. In the uv range, there is a relatively small absorption peaking near 390 nm. In Figure 3 absorption decays vs. time have been plotted for convenience on a semilog scale. The decay of the solvated electron as recorded a t 900 nm lasts 10 ps and apparently obeys no simple kinetics. From Figure 3 the half-time tll2, Le., the time at which the optical density equals half of the initial optical density is 570 ns. In spite of experimental errors and assuming the decay to be due to homogeneous concurrent pseudo-first- and second-order reactions, numerical calculations based upon nearly equal initial concentrations of the solvated electron and reacting species for the second-order reaction yield k(pseudo first) = 2 X lo5 s-l and k(second) = 1.75 X loll M-l s-l. At 400 nm the absorbing species decays more slowly than does the solvated electron except within the first microsecond during which the decay seems to parallel that of es-. The Journal of Physical ChemistryzVol. 80, No. 15, 1976
1688
Delaire et al.
TABLE I: Rate Constants and Yields in Hydrazinium Chloride and in BiDhenvl Solutions
I
I
0 570
[Phz], M tllz, 10-9 kz(Phz), lo7s-l kz, lolo M-l s-l G(Phz-), anions/100 eV
0 570
0.5
0.8
90 1.57
70 9.9 1.23
0.16 110
22
7.85
2.0 42
16.8 0.84
1.6
3.40 14
3.16 4.95 1.97 1.45 2.0 3.0 3.4 a t 112 is the time at which the optical density equals half of the initial optical density.
-1
0.051
[NzHtiClI, M tllz,a 10-9 kl[NzHs+],IO6s-l kl, 107 M-I s-l
i
WAVELENGTH [nm)
Figure 1. UV visible spectra of the pure hydrazine at the end of the pulse (0)and 4 ps after the pulse (A);2 M N2H5CIsolutions, 400 ns after the pulse ( 0 ) .
0.1
0.01
i
2
b
5
8
1
8
TIME (ps)
Figure 3. Decay of transients in pure hydrazine at 900 nm (0)and 400 nm (A).Decay in 2 M N,H& solutions at 900 nm ( 8 )and 400 nm
(A).
I 200
I
"
W
I
I
400
I
I
BOO
800
1000
WAVELENGTH (nm)
Figure 2. Transient spectra of pure liquid hydrazine at the end of the
1.6 X pulse ( 6 refers to the solvated electron) 0 ; M biphenyl solution, 600 ns after the pulse (partial scavenging of e=-) 0 . Inserts represent 1.6 X M biphenyl solution (a) h 1000 nm, 200 ns/division; (b) h 400 nm, 100 ns/division.
Concentrationshigher than 0.5 M hydrazinium chloride are necessary to obtain a change of kinetics. In the presence of acidic ions NzH5+, the decay of the solvated electron follows a pseudo-first-order law due to the reaction e,-
+ NzH5+
-
NzH3 + Hz
(1)
Table I shows the concentration dependence of the rate constant due to the effect of ionic strength. Because of the high value of the dielectric constant of hydrazine ( E = 53 at 20 "C), the Debye-Huckel theory17applies and the dependence of the rate constant h , on ionic strength I reads log k l = log k10 -
2Bv7 + 2 C I 1 a a 4 where k10 is the rate constant at zero ionic strength, a = 4.5 X cm the distance of closest approach, C a semiempirical parameter, B = 0.935 M-lIZ, and a = 0.403 X lo8 M-l12 cm-l.
+
The Journal of Physical Chemistry, Vol. 80,No. 15, 1976
The data of Table I on the decrease of k l vs. hydrazinium chloride concentration show, in contrast to conductimetric results,15 that C may be neglected. Then numerical calculations give h1° = (5 f 3) X 107 M-1 s-1. Figure 1shows the transient absorption spectrum in a 2 M NzH5Cl solution 400 ns after the pulse; that is at a time longer than the electron decay. It should be noted that there is an absorption band from 330 to 450 nm which peaks at 370 nm. There is no evidence for identifying the absorbingspecies with that observed in pure hydrazine except the fact that both species decay within similar kinetics above the first microsecond as shown in Figure 2. Therefore we propose that the uv absorption would be partially due to the tail of the solvated electron spectrum (rapid decay) and partially due to an oxidizing species (slow decay). 2. Yield of the Solvated Electron. In contrast to liquid ammonia, alkaline metal solutions in hydrazine are unstable on the time scale of seconds and explosive,l* so that there is no possibility to use them to determine the extinction coefficient of the solvated electron and, hence, the electron yield by direct measurement. Therefore biphenyl (Phz) was used as electron scavenger in order to obtain the G value of the solvated electron. At the end of a 3-ns pulse the transient absorption spectrum in a 1.6 X M biphenyl solution in hydrazine is identical with that obtained in pure hydrazine. No fluorescence emission has been detected in these experiments. In Figure 2 the 600-11s transient spectrum shows two bands ' peaking at 400 and at 640 nm and a broad absorption beyond 950 nm. This spectrum has been assigned to the biphenyl anion because of the similarity with the bands obtained from chemically prepared biphenylide anion19pz0and with those observed from pulsed biphenyl solution in cyclohexane21.22
1689
Nanosecond Pulse Radiolysis of Hydrazine and in ~ a t e r . ~ The 3 fact that there is no absorption near 950 nm at the highest biphenyl concentration excludes the presence of the solvated electron. Moreover as the absorption beyond 950 nm increases with the biphenyl concentration, the ir absorption cannot be due to the solvated electron in equilibrium with the biphenylide anion as recently pointed out in tetramethylsilane.24 Therefore the absorption beyond 950 nm has been assigned to the biphenylide anion. Figure 2 (inserts) shows typical traces of the absorption recorded respectively at 400 and at 1000 nm in a 1.6 X 10-4 M biphenyl solution. At 400 nm the initial absorbance reaches a maximum over a time corresponding to the disappearance of the absorptian at 1000 nm. Therefore the growth of the biphenylide anion is correlated to the decay of the solvated electron accordingly to the reaction e,-
-+ Phz
--*
Ph2-
(2)
Assuming no decay of biphenylide anion during the buildup because its lifetime is comparatively long (10 ps), the maximum of absorbance corresponds to solvated electrons scavenged at a specified biphenyl concentration. Table I gives the concentration dependence of the observed pseudo-first-order rate constant kz[Phz] and of the G(Phz-) value. The observed rate constant was calculated from the electron decay at 1000 nm and from the growth of biphenylide anion at 400 nm. The G (Phz-) value was obtained from the maximum absorbance recorded at 640 nm and is based on an extinction coefficient t(Phz-) = 1.21 X lo4 M-l cm-l at 640 nm, the mean value from ref 19 and 25. A t the highest biphenyl concentration, the half-time of the solvated electron is 14 ns. Compared to the pulse duration (3 ns) and to the half-time of the solvated electron in pure hydrazine (570 ns) it can be assumed that all solvated electrons present a t the end of the pulse are scavenged by 3.4 X lom3M biphenyl. Therefore the value G(Ph2-) = 3.4 f 0.2 obtained at the highest concentration can be considered as close to the yield Ges- in pure hydrazine in the nanosecond time range. Therefore the extinction coefficient €(es-) = (1.85 f 0.10) X lo4 M-l cm-l a t 1000 nm has been calculated from the experimental value of the product G X t. From the observed pseudo-first-orderrate constant the specific second-order rate constant of the scavenging has been found: kz = (1.7 f 0.3) X 1O1O M-l s-l. Discussion The observed absorption spectrum of the solvated electron in pure hydrazine has no maximum in the wavelength range 450-1100 nm. As previously pointed outll such an electron absorption spectrum would be probably due to electrons interacting more weakly with surrounding hydrazine molecules than in the case of water or alcohols in spite of the fact that the hydrazine dielectric constant ( E = 53) is between that of water and those of alcohols. However assuming the reaction of the solvated electron with biphenyl to be diffusion controlled, the comparatively low value of the specific rate constant kz = 1.7 X 1O1O M-l s-l corresponds to a diffusion coefficient of the solvated electron similar to that in water26 or in liquid a m m ~ n i a . Therefore ~' the maximum of the spectrum would be expected to be located in the same wavelength range as in ammonia or in amines.28 As concerns the decay of the solvated electron in pure hydrazine, some mechanisms can be ruled out. Indeed, in spite of the exploding character of alkali metal solutions in hydrazine,lS the blue color due to the solvated electron (possibly associated with metal) can be observed during a few seconds:
this implies very low rate constants for the reactions of solvated electron with another electron or with hydrazine. This fact incidentally excludes the reaction of the electron with the solvent previously suggested to explain the first-order kinetics of the electron decay observed in early pulse radiolysis experiments.11 On the other hand the reaction of the solvated electron with the acidic cation N2HbCcannot account for the electron decay because the rate constant k l is too low as is the corresponding rate constant of the hydrated electr0n.~9 Contrastingly H atoms generated either by excitation processes or by ion recombination promptly react with hydrazine29+30 and then they cannot be involved in the electron recombinations. Then it follows that the solvated electron will disappear by reacting with oxidizing species, probably NzH3.. As concerns the electron yield related to a 3-11s pulse, the G values in hydrazine (3.4) and in ammonia (3.0) were calculated relatively to the G value in water (3.3) so that our results are self-consistent. Under the circumstances, it seems that hydrazine is a singular case for which the electron yield is slightly higher than that in water. Similar to liquid ammonia, the value G = 3.4 can be reasonably explained by the slowness of the electron-NzH5+ recombination, which increases the escape probability of the solvated electron. Supposing an initial electron yield near 5, we have to ask ourselves which reactions can lead to a yield of 3.4 at 3 ns. Recent picosecond results in water31@ clearly demonstrate the occurrence of spur reactions between electron, H30+, and OH., thus explaining a decrease in the yield of electron from 4.8 to 3 in the nanosecond time range. In hydrazine, once the electron is solvated, its capture by the cation or by another electron is excluded as shown in this work. We mentioned above the scavenging by radicals being responsible for the decay after the pulse. This reaction is already effective at early times, especially as the primary species are concentrated in the spurs, and therefore lowers the escape probability of the solvated electron. Hydrazine is thus an example illustrating the faEt that chemical reactivity of ions must be taken into consideration as well as physical properties of the solvent. Consequently the knowledge of rate constants of the electron-cation neutralization would enable us to appreciate the extent to which present theoretical models are adapted to liquids.
Acknowledgment. The authors wish to thank Drs. J. C. Goudeau and M. Broussely for valuable discussions. References and Notes (1) 6. Ward, Adv. Chem. Ser., No. 81, 601 (1968). (2) J. Bellonl, P. Cordier, and J. Delaire, Chem. Phys. Len., 27, 241 (1974); J. Belioni and E. Saito in "Electrons In Fluids", J. Jortner and N. R. Kestner, Ed., Springer, Berlin, 1973, p 461. (3) N. A. Seddon, J. W. Fletcher, F. C. Sopchyshyn, and J. Jevcak, Can. J. Chem., 52, 3269 (1974). (4) Farhataziz, L. M. Perkey, and R . R. Hentz, J. Chem. Phys., 60, 717 (1974). (5) J. M. Brooks and R. R. Dew'eld, J. Phys. Chem., 75,986 (1971). (6)Farhatazlz, L. M. Perkey, and R. R. Hentz, J. Chem. Phys., 60, 4383 (1974). (7) L. F. Audrleth and B. A. Ogg, "The Chemistry of Hydrazine", Wiley, New York, N.Y., 1951. (8)H. Luclen and M. L. Plnns, NASA Tech. Note, TN D.2452 (1962). (9) U. PrBsch, 2.Chem., 4, 395 (1964). (10) W. E. Shelberg, USNRDL-TR 1002, 1966. (1 1) N. A. Slavinskaya, B. M. Kozlov, S.Va Pshezhetskil, F. V. Shemarov, and L. V. Tshepel, Proc. Tlhany Symp. Radlat. Chem., 3rd, 1077, 1219 (1972); N. A. Slavlnskaya, B. M. Kozlov, N. V. Petrukhln, S. A. Kamenetskaya, and S.Ya. Pshezhetskil, Khim. Vys. Energ., 8, 68 (1974). (12) P. Cordier and J. Delalre, to be submitted for publicatlon. (13) J. Delaire, These 38 cycle, Orsay, 1973. (14) E. Salt0 and J. Bellonl, Rev. Scl. hstrum., submitted for publlcatlon. The Journal of Physlcal Chemlstry, Vol. 80, No. 15, 1078
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J. Pola, G. Levin, and M. Szwarc
(15) M. Brousseiy, These doctorat, Poitiers, 1975. (16) J. K. Thomas and R. V. Bensasson, J. Chem. Phys., 46,4147 (1967). (17) P. Debye and E. Huckei, Phys. 2..24, 185 (1923); E. A. Guggenheim, Phil. I%& 19, 588 (1935). (18) P. Pascal, "Nouveau Trait6 de Chimie Minerale", Masson et Cie, Paris, 1956, tome X, p 560; T. W. B. Welsh and H. J. Brcderson, J. Am. Chem. SOC., 37, 816 (1915). (19) K. H. Buschow, J. Dieleman, and G. I. Hoijtinck, Mol. Phys., 7, 1 (1963). (20) G. I. Hoijtinck, Chem. Phys. Lett, 26, 318 (1974). (21) J. P. Keene, E. J. Land, and A. J. Swallow, J. Am. Chem. Soc., 87, 5284 (1965). (22) J. K. Thomas, K. Johnson, T. Kiippert, and R. Lowers, J. Chem. Phys., 48, 1608 (1968). (23) J. H.Fendler, H. A. Giilis, and N. V. Kiassen, J. Chem. SOC.,Faraday Trans. 1, 70, 145 (1974).
(24) J. M. Warman, M. P. de Haas, E. Zador, and A. Hummel, Chem. Phys. Left., 35, 383 (1975). (25) D. Gill, J. Jagur-Orodzinski,and M. Szwarc, Trans. Faraday SOC.,60, 1424 (1964). (26) K. L. Schmidt and W. L. Buck, Science, 151, 70 (1966). (27) E. C. Evers and P. W. Frank, J. Chem. Phys., 30, 61 (1959). (28) L. M. Dorfman and F. Y. Jou. "Electrons in Fluids", J. Jortner and N. R. Kestner, Ed.. Springer, Berlin, 1973, p 447. (29) J. Beiloni and M. Haissinsky, lnt. J. Radiat. Phys. Chem., 1,519 (1969); E. Hayon and M. Simic, J. Am. Chem. SOC.,94,42 (1972). (30) M. Schiavello and G. G. Voipi, J. Chem. Phys., 37, 1510 (1962). (31) R. K. Wolff, J. E. Aldrich, T. L. Penner, and J. W. Hunt, J. Phys. Chem., 79, 210 (1975). (32) C. D. Jonah, M. S. Matheson, J. R. Miller, and E. J. Hart, to be submitted for publication.
Equilibrium and Kinetic Studies of Disproportionation of Sodium Tetracenide in Benzene. The Effect of Added Tetrahydrofuran J. Pola, G. Levin, and M. Stwarc* Department of Chemistry, State University of New York, College of Environmental Science and Forestry, Syracuse, New York 13210 (Received February 17, 1976) Publication costs assisted by the National Science Foundation
The equilibrium and the kinetics of disproportionation of sodium tetracenide (Te.-,Na+) in benzene containing small amounts of tetrahydrofuran (THF) was investigated. It was shown that the equilibrium is represented by the stoichiometric equation, 2Te.-,Na+,(THF), @ Te Te2-,2Na+,(THF)2,-2 2THF and the formal equilibrium constant Kdispr = [Te][Te2-,2Na+]/[Te.-,Na+l2 varies from 400 at very low concentration of THF to in bulk THF. The kinetics of disproportionation was investigated by flash-photolytic technique leading to the value of k-dispr of 1.5 X lo9 M-l s-l for the reaction Te2-,2Na+,(THF)zn-2 Te 2Te--,Na+, (THF),- 1.
+
+
+
It is generally believed that the disproportionation of planar aromatic radical anions, a reaction that yields dianions and the neutral hydrocarbons, is highly endothermic and its equilibrium constants are very l ~ w . l Examples -~ to the contrary are but all of them refer to systems in which the geometry of a t least some of the reacting species is drastically changed; this allows us to understand the preference in these systems for dianions over the radical anions. It was surprising, therefore, to find out7that the disproportionation of the planar tetracenide radical anions (Te--,Cat+), derived from the planar hydrocarbon, is favored in diethyl ether, e.g., Kdispr of lithium tetracenide in this solvent exceeds 1 2Te--,Cat+
* Te + Te2-,2Cat+
Kdispr
In fact, as shown by the data collected in Table I, the disproportionation equilibrium of Te.-,Cat+ is strongly affected by the nature of cation, and even more by the nature of solvent for reasons discussed elsewhere7 (see also ref 2 and 3). We wish now to report that the disproportionation constant of sodium tetracenide is even larger in benzene and its value exceeds 400. Moreover, we wish also to describe the spectacular changes in the value of Kdispr in benzene resulting from the addition of small amounts of tetrahydrofuran (THF).
Results Tetracene in THF solution was reduced on a sodium mirror. The reduction was not carried out to completion and the The Journal of Physical Chemistry, Voi. SO, No. 15, 1976
-
concentrations of the radical anions and of the unreduced hydrocarbon in the resulting solution were determined spectrophotometrically. Thereafter, THF was distilled off and an equal volume of benzene was distilled in. All these operations were performed on a high-vacuum line. The change of solvent led to a dramatic result. The concentration of the hydrocarbon in the benzene solvent substantially increased (the increase corresponded to about one-half of the original concentration of Te--,Na+). For example, a prepared THF solution was 1.6 X 10-5 M with respect to Te and 6.1 X 10-5 M with respect to Te.-,Na+. After replacement of THF by benzene the concentration of Te increased to 4.8 X M, while the conversion of 2Te--,Na+ to Te Te2-,2Na+ would make the concentration of tetracene equal to 11.6 (0.5)6.1] X 10-5 = 4.65 X M. This implies that the reaction
+
+
2Te.-,Na+
-
Te
+ Te2-,2Na+
converts nearly all the Te.-,Na+ into Te2-,2Na+. Due to its low solubility, only a fraction of the formed Te2-,2Na+ was dissolved, its concentration was found to be 0.22 X M. However, in spite of the large excess of Te, no Te.-,Na+ was detected in the investigated solution, while a concentration -5 X 10-7 M still could be measured. From such spectrophotometric data we conclude that the disproportionation constant in benzene is greater than 400. Reduction of tetracene in benzene containing small
