Bradley Bockrath and Leon M . Dorfman
2618
Kinetics of Electron Transfer from Aromatic Radical Anions to Alkyl Halides in Tetrahydrofuran. Effects of Sodium Cation Pairing' Bradley Bockrath and Leon M. Dorfman* Department of Chemistry, The Ohio Sfafe University, Columbus, Ohio 43270
(Received June 29, 7973)
Pubhation costs assisted by the U . S. Atomic Energy Commission
Absolute rate constants have been determined, by the pulse radiolysis method, for electron transfer from naphthalenide and biphenylide ions to n-butyl bromide and n-butyl iodide in tetrahydrofuran solution. The sodium cation-paired form of the radical anion, (Na+. A. -), exhibits reactivity that is roughly two orders of magnitude lower than that of the free-radical anion, A * - . The temperature coefficient for the reaction was interpreted in terms of an equilibrium of loose and tight ion pairs, the former exhibiting greater reactivity. A small effect of added salts upon the rate constants was observed.
Introduction
We have determined, by the pulse radiolysis method, the effect of sodium cation pairing on the rate constants for electron transfer from aromatic radical anions as donors to alkyl halides in tetrahydrofuran solution. Effects of metal cation pairing on the reactivity of aromatic radical anions in several types of reaction have been establishedZ-? earlier, although limited data pertaining to electron transfer are available.8 The pulse radiolysis method, which has been used extensively in the study of electron transfer kinetic^,^^^^ provides a convenient means of studying such cation-pairing effects since the reaction may be observed with metal cations present or absent, as one chooses. Pulse radiolysis of solutions of aromatic molecules in liquids such as tetrahydrofuran yields the radical aniong-ll by attachment of the solvated electron e,-
+
A = A.-
(1)
In the presence of free excess Na?, the ion pair1' with the solvated electron, (h'a+, es-), is formed (hz = 7.9 X 10l1 M - l sec-1 in THF) e,-
+
~
a = + ("a+,e s - )
(2)
This species will also attach, with a high rate constant,ll to aromatic hydrocarbons to yield the ion-paired radical anion (Na', e 3 - )
+
A = (Na', A , - )
(3)
Observation of either A m - or (Na+, Aa-), by fast optical absorption spectrophotometry, with nanosecond time resolution when necessary, permits comparative kinetic studies of the two species to be easily made. We have obtained such data for naphthalene and biphenyl radical anions reacting with n-butyl bromide and n-butyl iodide in T H F A , - + n-C,H,X = A + n-C,H,. + X(44 (Na', A.-) + n-C,H,X = A + n-C,H,. + Na'X(4b) Experimental Section
The source of the electron pulse, as in our earlier studies.lZ was a Varian V-7715A electron linear accelerator, delivering 3-4-MeV electrons a t a pulse current of about 350 mA for pulse duration of 100-1500 nsec and about 600 mA for pulse duration less than 80 nsec. Electron pulses The Journal of Physicai Chemistry, Vol. 77. No. 22, 1973
for 20-80-nsec duration were used in this work. The transient optical absorptions were observed using an H.T.V. 196 detector with an S-1 response. A Bausch and Lomb grating monochromator, type 33-86-25, f l 3 . 5 was used. Corning filters were used to eliminate second-order components from the analyzing light beam. Our standard reaction cells,lZ with high-purity silica windows and a cell length of 20.0 mm, were used in most runs with a double pass of the analyzing light beam. Accordingly the optical path length was 40.0 mm. Temperature coefficients of the rate constants were obtained with a thermostatic box heated or cooled by a stream of nitrogen gas. The T H F was purified first by refluxing under argon, for several hours, a solution containing benzophenone and excess sodium metal. This was followed by distillation through a glass bead-packed column, the middle fraction being retained. It was then degassed and vacuum transferred into a bulb containing a mirror of freshly triple-distilled potassium. I t was vacuum transferred from this bulb into the reaction cells just prior to the runs. The aromatic compounds used were zone refined, commercially supplied, with a nominal purity of a t least 99.9%. n-Butyl bromide (Fisher reagent grade) and nbutyl iodide (Matheson Coleman and Bell) were purified by washing with several portions of concentrated HzS04, followed by several portions of 10% sodium carbonate, and finally distilled water. The washed materials were dried and distilled through a glass bead-packed column, and the middle fraction was retained. n-Butyl iodide was stored in the dark over mercury. Sodium tetraphenylboron (Fisher reagent grade) was recrystallized as recommended,l3 and stored in uacuo until used. Sodium iodide (either Baker analyzed grade, 99.0870, or Alfa Inorganics optronic grade) and sodium perchlorate (G. Fredrick Smith reagent) were used without further purification. Results and Discussion
The reactions were observed by monitoring the decay of the aromatic radical anion a t the maximum of the long wavelength absorption bands. Naphthalenide was monitored a t 775 nm, bipheny!ide a t 630 nm. The differential rate equation for the observed process, reaction 4a or 4b, is -d[A*-]/dt = K,[A.-][n-C,H,X] (5)
Kinetics of Electron Transfer from Aromatic Radical Anions Under our experimental conditions, with [n-C4HgX]o >> [A.-]o or [n-C4HgX]o >> [(Na+, A.-)]o, the reaction follows a first-order rate law, and h g is readily obtained by linearization of the data in accord with the integrated form of ( 5 ) . In a typical series of experiments, the decay of A * - or ( N a + , A * - ) was monitored in the absence and in the presence of alkyl halide. In the absence of alkyl halide the decay of A . - followed second-order kinetics, while that of (Na+, A - - ) followed either second-order or, in some cases, mixed first- and second-order kinetics. With alkyl halide present the decay was shown to follow a first-order rate law over a range of alkyl halide concentrations of about one order of magnitude. The initial condition [nC4H9X]o/[A.-]o was such that the half-life of the radical anion in the absence of n-CdHgX was a t least ten times its half-life in the presence of the electron acceptor. The pseudo-first-order rate constants, hf4a or k’4b, obtained from first-order plots under these conditions, were then plotted against [n-C4HgX] to determine kca or k4b. The data exhibited no dependency of the rates on the concentration of the aromatic molecule, indicating that there was no measurable back reaction. The kinetic order in butyl halide was unity, the slopes from plots of log k4 us. log [n-C4HgX] falling in the range 0.94-1.08. As indicated, the rate of reaction of the donor anion with the solvent counterion was, in most cases, negligible relative to the rate of reaction 4. Where this was not the case it was taken into account in determining h4. In the case of reactions involving biphenylide ion as donor, the following complication necessitated a modification in the method of determining h4 from the observed rate curves. In the biphenylide systems the absorption a t 630 nm did not decay to zero, but reached, instead, a weak residual plateau after the initial decay. At 630 nm the residual plateau was typically about 5% of the absorption measured a t time zero. Time resolved absorption spectra, Figure 1, taken a t time zero and a t the plateau region, show the absorption remaining after the decay of the aromatic radical anion. Several species might reasonably contribute t o the residual absorption. The biphenyl radical cation,l4 the biphenyl triplet,lS and the phenylcyclohexadienyl radical16 all have uv absorptions that could contribute to a composite spectrum. The long wavelength component may possibly be due to the radical cation, while the short wavelength component could be due to any one or a mixture of the three species. Where such a plateau is formed, the pseudo-first-order rate constants, kf4a, were obtained from plots of In (Dt - D m ) us. time. This method is valid if the residual absorption, D,, does not change during the course of the reaction with butyl halide. The D , plateau shows negligible decay on the time scale of observation of reaction 4a. The possibility that the species responsible for the D , absorption grows in during reaction 4a was also ruled out because plots of In (Dt - D a )were always linear and values of k4, were constant over a wide range of butyl halide concentration. As a further check, values of h4a obtained at 400 nm, where the residual absorption constitutes a higher percentage of the initial absorption, were found to be equal to those determined a t 630 nm. Furthermore, there is no growth of absorption a t these wavelengths after the end of a pulse in solutions of biphenyl without added butyl halide. The formation of the D , absorption is therefore most likely complete by the end of pulse.
2619
4.06
i m 4
.04 -
.0201
1
I
35 0
I
450
5 50 65 0 WAVELENGTH ( nanometers)
750
850
Figure 1. Spectra of THF solution 5.5 X M in biphenyl and 2.2 X I O p 3 M in n-butyl iodide showing absorption due to biphenylide and the residual absorption plateau after reaction 4a is complete. Curve A taken at time zero following an 80-nsec electron pulse. Curve B taken 300 nsec following the electron pulse at the plateau of residual absorption. TABLE I : Rate Constants for Reaction 4a with n-Butyl iodide and n-Butyl Bromide in THF Rate constant, M - ’ sec-’ at 298°K
Biphenylide Naphthalenide
n-Butyl iodidea
n-Butvl bromidea
(9.6f 1.4)x 109 (7.4& 1.1)X IO9
(3.4 0.6) x 107 (3.3 f 0.6) X l o 7
*
aThe range of concentration used in determining kda was 5 X M for n-butyl iodide and 0.02-0.2 M for n-butyl bromide.
3X
It was necessary to use In ( D t - Dm)plots only for the reactions of free biphenylide ion. In other cases, D, is a negligible fraction of the initial absorption. In the former cases, a relatively high biphenyl concentration (0.06 to 0.3 M ) was used in order to overcome the competition for e,by the added butyl halide, whose electron-scavenging rate constants are very high.ll This would undoubtedly favor a higher fractional yield of the intermediates responsible for the D, absorption. R a t e Constants for Free Aromatic Radical Anion Values for k4a with both n-butyl iodide and n-butyl bromide as acceptor are shown in Table I. It is interesting to note that, whereas the electron transfer to n-butyl iodide is perhaps only slightly lower than the diffusion-controlled limit, the reaction with n-butyl bromide is some two orders of magnitude slower than diffusion controlled. Earlier studieslo of electron transfer from aromatic radical anions to a different neutral aromatic molecule. in solvents such as 2-propanol and ethylenediamine, had shown that the rate constants approached the diffusion-controlled limit when the reduction potential of the donor exceeded that of the acceptor by roughly 0.2 V. The results for the reaction with n-butyl bromide, which very likely involves a dissociative attachment. implies qualitatively that n-butyl bromide is more difficult t o reduce than either biphenyl or naphthalene.
Rate Constants for the Cation-Paired Radical Anion Values for h4b were determined by monitoring the decay of (Ka+, A * - ) for naphthalenide and biphenylide in the presence of butyl halide in a T H F solution containing sufficient free sodium ion to maintain the aromatic The Journal of Physical Chemistry, Vol. 77, NO. 22, 1973
Bradley Bockrath and Leon M. Dorfman
2620
radical anion in the cation-paired form. Th.- was achieved by addition of dissociative sodium salts to the solution. Three salts were used in separate experiments: sodium tetraphenylboron (for which the dissociation constant in T H F is knownl3), sodium iodide, and sodium perchlorate. The concentration of sodium ion was well in excess of the amount required to maintain the equilibrium17 (Na',
A.-)
z=Z X a '
+
A
(6)
overwhelmingly in the cation-paired form. The decay of the ion-paired aromatic radical anion was monitored in the absence as well as the presence of butyl halide. In the absence of butyl halide the decay was of second order or under some conditions of mixed order, and presumably consisted mainly of reaction with the solvent counterion. This rate of decay was roughly an order of magnitude slower than in the case of the free aromatic radical anions, indicating a counterion recombination rate constant of about 1010 M-1 sec-l for the cation-paired M - 1 sec-l for the free aromatic radical anion and radical anion. This seems reasonable since the counterion recombination for (Na+, A a - ) is an ion-dipole reaction rather than an ion-ion reaction. The values obtained for h4b in T H F solutions containF sodium tetraphenylboron are shown in ing 1-2 x Table II. Comparison of these values with the corresponding data in Table I reveals that the reactivity of the sodium cation paired aromatic ion in these electron transfers is fully one to two orders of magnitude lower than the reactivity of the free anion. I t is also interesting to compare our values of k4b with the rate constant of 4.3 x M - 1 sec-1 measured by Garst and BartonlS for the reaction of sodium naphthalenide with either hexyl or octyl fluoride in DME. The reactivity of the alkyl halides thus spans more than 12 orders of magnitude over the series from iodide to fluoride. In presenting the data in Table 11, it was essential to cite the particular added salt used, as well as its concentration, for the following reason. The value of the rate constant. as it turns out, is dependent upon both the choice of added salt and its concentration, showing an increase with increasing concentration. This may be seen very clearly in Figure 2 which shows the vaiue of the observed rate constant for the electron transfer from sodium biphenylide to n-butyl iodide as a function of concentration for all three added salts. The increase in the observed rate constant with increase in added salt concentration is not due to reaction of the donor radical anion with the added salt. This was shown by determining the lifetime of (Na+, A - - ) in the absence of butyl halide, a t an increased concentration of added salt, It was found that the rate of decay of sodium biphenylide increased to the extent of only 50% for an increase in sodium perchlorate concentration from 0.02,5 to 0.31 F. Since the half-time for this decay is always more than ten times the half-time in the presence of butyl iodide (at all concentrations used), direct reaction of (Na+, A -) with the perchlorate, if i t occurs a t all, cannot be the cause of the increased rate constant in Figure 2. The concentration of added salt does, however, affect the value of k4b, as may be seen in Figure 2 which shows a nonlinear increase in the observed rate constant with increase in salt concentration. The sensitivity to saIt coiicentration depends strongly on the nature of the anion of the salt, perchlorate showing the largest effect and tetraphenylboron a very small effect. The solid curves in this The Journal of Physicai Chemistry, Voi. 77, No. 22, !573
4
i
I 0
I
I
0.04 0.08 0.12 SALT CONCENTRATION (FORMAL)
Figure 2. Effect of added salts on h4b. Solid lines pertain to concentrations of added sodium salts sufficient to assure biphenylide is completely associated with sodium cation. Dashed lines represent extrapolation of k4b to infinite dilution of added sali. The rate constant for reaction of free biphenylide, kqa, has been added for comparison: A , sodium perchlorate; 0 ,sodium iodide; 0 , sodium tetraphenylboron; 3 , k4a for reaction of free biphenylide without added salt. TABLE JI: Rate Constants for Reaction 4b with n-Butyl Iodide and n-Butyl Bromide in THF Solution Containing 1-2 X l o - * F of Sodium Tetraphenylboron
-
Rate constant, M - ' sec-', at 298'K
Na+, biphenylide Na+, naphthalenide
n-Butyl iodidea
n-Butyl bromide*
(4.3f 0.2) X IO8 (9.3 f 0.5) X l o 7
(1.3 f 0.1) X lo6
