Pulse Radiolysis of Liquid
Ammonia
The Journal of Physical Chemistty, Vol. 82,No. 5, 1978 537
Primary Processes Studied by Pulse Radiolysis of Liquid Ammonia. 2. Influence of a Chemically Inert Electrolyte J. Bellonl," P. Corder, J. A. Delalre, and M. 0. Delcourt Laboratoire de Physico-Chirnie des Rayonnernents, Associ6 au CNRS, Universit6 de Paris-Sud, Centre d' Orsay, 9 1405 Orsay, France (Received July 27, 1977) Publication costs assisted by Laboratoire de Physlco Chimie des Rayonnements, Orsay
The main reaction of solvated electrons in irradiated liquid ammonia appears to be with the oxidizing radicals NH2and NH although the product, the NH, ion, is not observed in the pure medium. Addition of a chemically inert salt such as NaC104makes this observation possible. Such a result, interpreted as a screening effect on the electrostatic attraction between the geminate pair, strongly supports the proposed hypothesis of ion pairing between the electron and a cation prior to reaction with radicals in the pure medium. This hypothesis is discussed in relation to some earlier results and to the present resuIts and permits an estimation of the distance between the geminate ions (85 A).
Introduction In the preceding paper' we elucidate the nature of the UV absorption bands induced by pulse radiolysis of liquid ammonia. We are now interested in the fate of the solvated electron in this medium. The relatively high value of the ammoniated electron yield has been interpreted as a result of the slow reaction of earn-with NH4+.2-5 Nevertheless, an important question remained about the fate of earn-;the product NH2- expected from its reaction with NH2 and NH radicals3s6has never been observed in the pure medium, at least in sufficient amounts to support the hypothesis of this reaction. A low intermediate concentration would have been interpreted in terms of a high rate constant for the reaction of NH2- with NH4+. In metalammonia solutions, though NH2- has been observed, its rate of neutralization with NH4+ has been found to be slow.' To explain this conflicting situation, we suggested1 that the electrostatic perturbation caused by the added electrolyte Na+ + em- had a determining importance, even at the low concentrations used. Thus, to distinguish chemical scavenging from the electrostatic effect, we studied solutions of chemically inert electrolytes. Experimental Section The pulse-radiolysis set up, the optical detection, and the procedure for preparing solutions have been described in part 1.l NaC104.6H20 was from BDH and NH4C1O4 from Prolabo. Results The chemically inert electrolyte used was NaC104. The C104- ion is well known for its nonreactivity toward e; or radicals produced radiolytically in water and also in liquid ammonia, both from y radiolysis3,' and pulse radiolysis results.6 We used NaC104 in the same concentration range as metal-ammonia solutions1 in order to obtain a comparable electrostatic effect. Contrary to what happens in the pure medium, and in the same way as in metal ammonia solutions, a build up of amide ions is observed after the pulse (Figure 1). Although the end-of-pulse spectrum (Figure 1, curve a) and the decay of earn-in the infrared are not appreciably altered with respect to the pure medium, the UV spectrum changes into a new one during the first 150 ns (Figure 1, curve b). The increase of absorbance near 330 nm is correlated with the decay of em- and, at 150 ns, the optical density of earn-is about half its initial value. 0022-3654/78/2082-0537$01 .OO/O
When subtracting the decay of both X and Y bands, supposed to be identical with that of earn-in the same solution, the true evolution in time of the new species is obtained (insert of Figure l ) , and its spectrum a t 150 ns (Figure 1,curve c) is very close to that of NH2- (Figure 1, curve d).9J0 The absorbance at 330 nm, measured 150 ns after the pulse, which corresponds to the maximum in time, is plotted vs. NaC104 concentration in Figure 2; above M there is an appreciable increase in optical density, up to a plateau value reached for M. The corresponding G(NHc) is then about 1.5, Le., half the value of the initial G, - and exactly the same as the yield of earn-consumed wi%n the same time. It is noteworthy that in these NaC104 solutions the risetime of N H c absorbance is 0 deduced considerably shorter (150 ns) than that ( ~ 4 0 ns) from homogeneous kinetics in Na solutions with h3 = 3.5 X 1O'O M-l s-l.' This means that the spatial distribution of radiolytic species earn-and NH2 is still nonhomogeneous in this time range. Beyond 150 ns, the decay of NH2- is assumed to be due to the recombination with NH4+. If the concentration of the cation is deduced from charge balance, a second-order rate constant is estimated at (3 f 1) X 1O1O M-l s-l close to the calculated diffusion-controlled rate constant (Table I). This value is a factor of 40 larger than the corresponding value in Na so1utions.l The difference is discussed below in relation to the electrostatic interactions between ions. The most important feature of the NaC104 effects lies, contrary to the results obtained in the pure medium, in the important yield of NH2- quite unexpected from classical ionic strength effects. Indeed such an influence of chemically inert electrolytes on the evolution of species has never been observed so far in radiolysis. It is therefore of interest to mention that the results of this work concern numerous solutions in separate runs. We also reinvestigated solutions of NH4C104 at concentrations up to 10-1 M; no noticeable difference in spectrum and kinetics (markedly in the UV range) was observed with respect to the pure solvent, at least during the first microseconds.
Discussion Neutralization Rate Constant Between NH; and NH4+. Reasons for the difference (Table I) of hNH2-+NH4+ in Na and NaC104 solutions must be found. First of all, the 0 1978 American Chemical Society
538
Belloni et al.
The Journal of Physical Chemistry, Val. 82, No, 5, 1978
TABLE I: Rate Constants in Liquid Ammonia at -50 "C Reaction ;e, NH,' earn + z" (NH,-, Na') + NH,' (NH,-, Na+) + NH,'
Ref
+
1.7 x 2.0 x 4.5 x 3.6 x
e
10-4 10-4 10-5 10-5
5.4
4.45 x 10"
4.1 4.5 6.3
1.7
lo6
3.5 x 1010 7.0 X l o 8 ( 3 * 1)x 1 0 ' 0
6.2 x 1O1O X
11-12 1 1 This work
1O'O
a Mutual diffusion coefficient from D(em-) = 1.5 x lo-, cmz s-l,13D(NH,+) = 2.3 X lo-' cmz s-' (from h o= 142 n-' cmz s-l (from the equiv-' cmz a t -33.5 'C14 and Walden's rule; D(NH,) = 5.10-5cm2 s-l, and D(NH,-, Na+)= 1.3 X Reaction radius (sum of radii Stokes-Einstein equation assuming R(NH,) = 1 A and R(NH,-, Na+)= 4 A , respectively). = 3.1 A (from molar volume = 70 mL1'); R(NH,+)= 2.3 A (Stokes-Einstein radius). Diffusionof reactants): R(e,-) controlled rate constant at zero ionic strength, calculated according to the Smoluchowski-Debye equation (6(NH,) = 23.5 M). The value has at - 50 'CL3). In metal-ammonia solutions ([Na] 3.2 X lo-, M). e In NaC10, solutions (J been deduced from the Figure 1 (insert, full line). The experimental error on kat includes an eventual formation of NH,slowly occurring in the 0.1-1-ps time range beyond the maximum observed.
*
Moreover the IR spectrum is not modified in the experiments where NH4C104is added to ammonia, albhough the pair has already been formed by a diffusion-controlled process. This implies the identity of the spectra of emand (eam-,NH4+). Reexamination of the Fundamental Processes in Pure Liquid Ammonia. Fate of the Solvated Electron, In contrast to solutions of electrolytes, the only ions present in irradiated neat ammonia are the species em- and NH4+. Concerning the formation and direct observation of NH, in this medium, it must be pointed out that the electron is strongly attracted by the cation. The competition between the diffusion of earn-toward NH4+and the diffusion of em- toward the neutral radical NH2 is then in favor of the reaction
*-
E
t 500 VA~LLffiTl4hm)
2%
3W
Figure 1. Transient absorption spectra In NaC104 solutions in liquid mol L-l): (curve a) ammonia at -50 OC ([NaCIO,] 2 5 X end-of-pulse spectrum; (curve b) spectrum 150 ns after the ulse (0 [NaCIO,] = 9.8 X IO-, mol L-l; 0, [NaCIO,] = 2.2 X 10-'mol L-l; A, [NaClO,] = 6.2 X IO3 mol L-I); (curve c) difference b - a12; (curve d) NH,- spectrum at -50 OC, normalized at 330 nm. Insert shows the evolution of absorbance at 330 nm in [NaC104] = 9.8 X IO4 mol L-': (dashed curve) experimental result; (dotted curve) absorbance in pure solvent; (full curve) difference between the two preceding ones (see text).
decays from which the rate constants were determined are too slow to be still disturbed by nonhomogeneous distribution. Secondly, concentrations of Na and NaC104 used are comparable and are in a such range that NH2is strongly paired with Na+ (Kd = 2.2 X M at -33 OC16), and probably reacts as a pair in both cases. The only possible difference therefore arises from the cation NH4+. The C10, anion is probably less associated with cations than is e -,the pairing of which with alkali cations is well known.lPIn fact, if NH4+forms ion pairs with added emin Na solutions, their lifetime must be long according to the low value of the rate constant kNpa++e,-,11,12so that NH4+ would react with (NH2-, Na+) in the paired form (em-, NH4+). The pairing could then markedly affect the reactivity of NH4+.ls Thus, the proposed mechanism is (i) in NaC104 solutions, amide ions are rapidly paired after their formation:' NH,
+ em-
NH,-
-+
Na'
c.(NH;,
Na')
(3)
Then the neutralization occurs with ammonium cations: (NH,-, Na') t NH,'
-f
2NH,
+ Na'
(8' 1
+
(ii) in Na+ em- solutions, reaction 3 is concomitant with the pairing of NH4+with added solvated electron: earn- + NH,'
(17)
T+ (eam-, NH,+)
These reactions are followed by (NH,-, Na')
+ (eam-, NH,')
+
2NH,
+ Na'
t earn-
(8")
-e,
+ NH,'
T+ (eam-,
NH,+)
(17)
This is true for both free and geminate ions since the secondary cation NH4+ and radical, resulting from an ion-molecule reaction, are closely spaced. Thus, the NH2 radical, when reacting with earn-,actually encounters the pair (eam-,NH4+) which has quit6 a low probability of disappearance by itself. In spurs, therefore, the diffusion of the electron toward the radical is assisted by the attractive field of the neighboring cation. The subsequent reaction of NH; with NH4+is then so fast that a stationary N H p concentration is never observed, even at long times, when spurs have diffused (at least at low temperatures): NH,
+ (eam-, NH,')
-+
(NH,-, NH,' )
-+
2NH,
(18)
This qualitative interpretation is also consistent with temperature effects in pure NH3. We observed that the decay of absorption at the maximum of the NH; spectrum near 340 nm is slowed down by raising the temperature from -50 to 0 0C.6 Perkey and Farhataziz'O observed a t 23 OC a slight increase in the absorption near 380 nm which lasts 60 ns. These features may be attributed to a small NH2- concentration, indicating a slowing down of the recombination NH2-, NH4+ with raising temperature; a possible explanation would be that ion pairing is weakened under such conditions. This interpretation is consistent with NH4C104results as well; the presence of the NH4+ ion can only enhance the formation of the pair (em+,NH4+) in a first step according to (17); then, due to mechanism 18, NH,- is not observed as a transient. Mechanism 18 proposed above (e;-radical reaction in the vicinity of the cation) to explain the lack of "2observation in the pure solvent is the only one which can explain the influence of NaC104. The effect of added electrolyte consists essentially in perturbing the specific attraction between the charges of the geminate pair. This screening effect makes the probability of encounter of e& with NH4+ or NH2 comparable so that the electron en-
Pulse Radlolysis of Liquid Ammonia
I
16
I I Illlll
104
I
The Journal of Physical Chemistry, Vol. 82, No. 5, 1978 539
I I111111
1
i3
L
I I I11111
[N$.
tos]’%not
account for certain anomalous r e s ~ l t s ; a~ pair, ~ * ~ stoi~ chiometrically equivalent and with spectral properties near those of the isolated electron, could be a more likely transient. In conclusion, in an irradiated liquid which has a low dielectric constant and where the main reaction of disappearance of e; is that with the parent radical as in ammonia, the perturbations introduced by the addition of a chemically inert electrolyte are observed through the formation of the resulting anion; the probability of first encounter by diffusion of the electron is in favor of the cation in the pure solvent, but becomes the same for the radical as for the cation in the presence of the electrolyte. This makes possible the observation of the anion as a distinct step.
I Ill
Flgure 2. Optical density at 330 nm and 150 ns after the pulse vs. concentration of NaCIO,. The dashed line expresses the optical density in pure liquid ammonia.
counters either the cation to form the long-lived pair (em-, NH4+)as in the pure medium, or the radical to form NHy outside the range of the NH4+ interaction (reaction 3). Because of this, the amide ion is observed during its diffusion toward a NH4+cation (Figure 1)in a stage which is distinct from the recombination (8’). Moreover it is remarkable that the limiting value of G(NHz-) (Figure 2) is just half that of the initial Gem-. Thus the perturbation introduced by the electrolyte provides a way to estimate the mean separation distance between geminate ions. In a single ion pair hypothesig, the probability p(r) dr that the nearest-neighbor solute ion (Na+ or C104-) will be formed in a spherical shell (r, r dr) around an electron may be expressed byzo p ( r ) dr = 4nr2n exp(4nr3n/3) (19)
+
The parameter n is the total number of ions per unit volume supposed randomly distributed. From relation 18, we can obtain the most probable distance P between an electron and the nearest neighbor ion:
r = (2nn)-”3
(20)
Now, if we insert the concentration of NaC104 corresponding to half the maximum effect (Figure 2), i.e,, 2 X M, we find P = 85 A. This value may be considered as an estimate of the mean distance between NH4+and esolv-in liquid ammonia at -50 “C. It is much higher than corresponding distances in ~ a t e r . ~ lIn- ~this ~ respect, ammonia resembles alcoholszzor hydrocarbonsz4although the relaxation time is very short as in water. With regard to the formation of a long-lived pair (eam-, “49, it would be of interest to reexamine some mechanisms for which a species such as NH4 was invoked to
Acknowledgment. This work was performed under contract with the Centre National de la Recherche Scientifique (ATP no. 2215). We gratefully acknowledge the technical assistance given by Mrs. Mifiana.
References and Notes J. Belloni, F. Billiau, P. Cordier, J. A. Delaire, and M. 0. Delcourt, J. Phys. Chem., preceding paper in this issue (reaction numbers in part 2 are the same as in part 1). J. Beiloni and E. Sako, Colloque Weyl I11 1972 in “Electrons in Fluids”, J. Jortner and N. R. Kestner, Ed., Springer, Berlin, 1973, p 461. J. Belloni, P. Cordier, and J. A. Deiaire, Chem. Phys. Lett., 27, 241
(1974). Farhatazlz, L. M. Perkey, and R. R. Hentz, J. Chem. Phys., 80,717
(1974). J. W. Fletcher, Discuss. Faraday SOC., 83, 18 (1977). J. Bellonl, F. Billiau, P. Cordier, J. A. Delaire, M. 0. Delcourt, and M. Magat, Discuss. Faraday SOC.,83, 58 (1977). J. Jove, Radiochem. Radioanal. Lett., 4,313 (1970). Farhataziz and P. Cordier, J . Phys. Chem., 80,2635 (1976). R. E. Cuthrell and J. J. Lagowski, Specfrosc. Lett., 1, 311 (1968). F. Billiau, these 36 cycle, Orsay, 1976. J. M. Brooks and R. R. Dewald, J. Phys. Chem., 75, 986 (1971). G. I.Khaikln, V. A. Zhigunov, and P. I . Dolin, Khirn. Vys. En., 5,
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(1975). S. Chandrasekhar, Rev. Mod. Phys., 15, l(1942-1943). H. Schwarz, J . Phys. Chern., 73, 1928 (1969). A. Mozumder, J . Chem. Phys., 50, 3153 (1969). W. G.Burns, R. May, G. V. Buxton, and G. S.Tough, Discuss. Faraday SOC.,83, 47 (1977)(and following discussions). N. Yamamoto, Y. Nakato, and H. Tsubomura, Bull. Chem. Soc. Jpn.,
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