J. Phys. Chem. le81, 85,1387-1391
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Spectroscopic Observation of Each Step in the Solvation of Ion Pairs. Li+N03- in Aprotic Solvents John P. Toth,+ Gary Rltzhaupt, and J. Paul Devlln’ Department of Chemistry, Oklahoma State Unlvers& Stlllwater, Oklahoma 74078 (Received: December 1, 1980; In Final Form: January 29. 198 1)
It is well established that the extent to which the degenerate antisymmetric stretching mode of NO3- is split into separate componentsis a sensitive measure of the strength of the interaction of the anion with an associated cation. It has also been demonstrated that the solvation of the cation, which can be divided into n stages in matrix experiments (nbeing the maximum solvation number), causes a discontinuous collapse of the u3(e)splitting. In this study advantage is taken of the relatively sharp character of the v3 component bands for each stage of solvation by aprotic solvents (ACN, THF, and Me2SO)to f i i y identify the v3 doublets for Li+NOp(solvent), in mixed argon-aprotic solvent matrices where m varies from 0 to n. In a form directly transferable to liquid solutions,such an assignment establishes the maximum solvation numbers for the contact ion pair [n = 4 (ACN), 3 (THF), 4 (Me2SO)]and the magnitude of the effect of each sequentially added solvent molecule on the cation-anion interaction.
Introduction The solvation of ion pairs by protic and aprotic solvents has been the subject of numerous experimental and theoretical studies. Primary objectives of such investigations have included (a) the determination of solvation numbers defined as the number of solvent molecules occupying the inner solvation sphere(s) of the ion(& (b) elucidation of both the nature and the strength of the solvent-ion interaction, and (c) the identification of the important species in solution, such as contact ion pairs, solvent-separated ion pairs, and dissociated ions, as well as the factors that affect the equilibria among these species. In general, information won through the direct study of bulk electrolytic solutions has been lacking in detail regarding these primary objectives. The complexities inherent in bulk solutions have led modern investigators to turn to theoretical methods1 and to experimental studies of simpler systems which yield direct insights to the behavior of molecular and ionic components of electrolytic solutions. In particular, Kebarle’s mass spectrometric investigations of the solvation of ions in the gas phase have yielded quantitative solvation information for dissociated ions2 while similar but perhaps less quantitative details regarding the solvation of contact ion pairs have been revealed through matrix-isolation infrared studies of the vapors over numerous high temperature halide3 and oxyanion salts.4 The matrix-isolation-contact-ion-pair work has, for the most part, been based on the fact that the splitting of oxyanion degenerate antisymmetric-stretching-mode frequencies by a polarizing cation charge is a sensitive and quantitative measure of the strength of the cation-anion interaction. Any factor which alters the strength of the interaction, such as the change from Li+ to K+ or the solvation of the Li+ cation, is readily noted by a sizeable change in the magnitude of the splitting, Aug, between the “degenerate” mode frequencies. For example, Av3 for Li+NO - varies from 262 cm-I for the isolated ion pair to 47 cm-B for the contact ion pair completely solvated in glassy NH3.4 For cases for which the geometry of the ion pair has been deduced, such as the bidentate structure for Department of Chemistry, University of Maine, Orono, ME 04473.
the Li+N03-species,lbVccorrelation curves are available to convert the magnitude of the observed Au3 values into quantitative values for the anion stretching force constants.& These values, combined with results for the same systems from ab initio Gaussian orbital calculationslbhave led to a considerable appreciation of the flow of electron charge during ion-pair solvation by protic molecules. For example, roughly one-half of the charge remaining on the Li+ ion in Li+N03-.H20, is canceled by addition of each successive water of hydration, while Av3 is reduced by approximately one-fourth by each hydration step. The quantitative data potentially available from the matrix isolation studies include the solvation number for contact ion pairs. For example, the introduction of increasing amounts of water or ammonia into primarily argon matrices can be used to reveal the stepwise collapse of Av3 that proceeds as the number of solvating molecules increases in a stepwise fashion.4c This stepwise collapse is particularly obvious in certain cases, such as the ammoniation of Tl+N03- for which a solvation number of five is indicated,4dbut, in general, precise hydration numbers have not been determined. This failure reflects primarily (a) a broadening and thus overlapping of the u3 infrared band components that accompanies hydration and (b) a significant distortion of the nitrate ion via interaction with water molecules such that Au3 is affected by direct anion-solvent interaction. Since H bonding is the likely source of the inhibiting factors, (a) and (b), it was proposed that similar matrixisolation studies of the various states of solvation, using aprotic solvents, would more clearly reveal the details of the solvation process.& With the details established for several simpler aprotic cases, additional confidence could (1) See, for example (a) E. Clementi and R. Barsotti, Chem. Phys. Lett., 69, 21 (1978);(b) J. C. Moore and J. P. Devlin, J. Chem. Phys., 68,826 (1978);(c) J. Amli5f and A. A. Ischenko, Chem. Phys. Lett., 61, 79 (1979). (2)See, for example (a) I. Dzidic and P. Kebarle, J. Phys. Chem., 74, 1466 (1970),(b) W.R. Davidson and P. Kebarle, J. Am. Chem. SOC., 98, 6125 (1976);(c) W.R. Davidson and P. Kebarle, ibid., 98,6133(1976). (3)B. S.Ault, J. Am. Chem. Soc., 100,2426 (1978). (4) (a) D. Smith, D. W. James, and J. P. Devlin, J. Chem. Phys., 54, 4437 (1971); (b) N.Smyrl and J. P. Devlin, J. Phys. Chem., 77, 3067 (1973); (c) G.Ritzhaupt and J. P. Devlin, ibid., 79, 2265 (1975);(d) G. Ritzhaupt and J. P. Devlin, ibid., 81,67(1977);(e) J. Toth, C. Thornton, and J. P. Devlin, J. Solution Chem., 7, 783 (1978).
0 1981 American Chemical Society
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The Journal of Physical Chemistty, Vol. 85, No. 10, 1981
' ' ' ' ' ' 100% ACN I
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Toth et al. I
I
r
I
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l
I l l I 100%THF
1
l
I
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100%Argon
I550
1450
1350
1250
Clli1
Flgure 1. Infrared bands for the v3 mode of LI+N03- isolated in matrices with compositions varying from pure argon to pure acetonitrile (ACN). The superimposed numerals identify v3 doublets with the number of solvating ACN molecules per ion pair. The ACN matrix concentrations in mole percent, corresponding to curves from bottom to top, are 0.0, 0.7, 5, 50, and 100.
then be brought to interpretation of the data for aqueous systems. The promise associated with this approach was enhanced by an infrared spectroscopic study of M+NO< pairs isolated in glassy pure matrices of acetonitrile (ACN), tetrahydrofuran (THF), and dimethyl sulfoxide (Me2SO), an investigation that (a) revealed relatively narrow bands for the solvated M+N03-, (b) confirmed that, as for H 2 0 and NH3matrices, the ion pairs remained contact in nature for such samples, and (c) clearly showed that the solvation in a glassy matrix very closely resembles that for the corresponding liquid solution.4e It has thus been clear that details of the solvation process that are obtainable from matrix-isolation studies would be directly applicable to liquid solutions of ACN, THF, and Me2S0. This paper reports such data for the single ion pair Li+N0< isolated in mixed matrices of argon with each of these three aprotic solvents with solvent concentrations varying from 0 to 100%. The choice of Li+NO, was dictated by the fact that, of the alkali metal nitrates, extensive aprotic solution data are available only for the lithium salt as well as by the greater range (and thus sensitivity) of Av3 values for lithium nitrate as compared to the other alkali metal nitrates.4a Experimental Section Matrix-isolated samples of Li+N03in either pure argon, pure aprotic solvent, or mixtures of argon with ACN, THF, or Me2S0 were prepared by vapor codeposition at 12 K onto a CsI substrate mounted within a vacuum shroud attached to an Air Products CS 202 closed-cycle helium refrigerator. The salt vapors were generated from anN
Flgure 2. Infrared bands for the v3 mode of Li+N03- isolated in matrices with compositions that vary from pure argon to pure THF. The superimposed numerals identify v3 doublets wlth the number of solvating THF molecules per ion pair. The THF matrix concentrationsin mole percent, corresponding to curves from bottom to top, are 0.0, 0.5, 12, 50, and 100.
hydrous reagent grade salts in pyrex Knudsen cells heated to the 400-500 "C range by using a nichrome resistance heater. The deuterated aprotic solventa, each having excellent transmission windows in the 7-pm region of the nitrate v3 mode, were obtained as deuterioquality reagents in sealed vials which were opened, and the content transferred to sample tubes, under dry nitrogen. After freeze-thaw degassing under vacuum the solvents were mixed to the desired ratio with argon prior to deposition. In the case of the relatively volatile solventa ACN and THF, a 2-L bulb was filled with the solvent gas at the solvent ambient vapor pressure and argon was subsequently bled into the bulb to achieve the desired ratio. The mixed gas was then metered through a glass flowmeter for codeposition with the salt vapor. In the case of the less volatile Me2S0, a dynamic mixing procedure wherein separate gaseous streams of argon and Me2S0 were brought to confluence immediately prior to the flowmeter gave the desired control of the argon:Me2S0 ratio. The spectroscopic data for the ACN and THF samples were obtained by using a Beckman IR-7 infrared spectrometer. Me2S0 sample spectra have been recorded by using a Digilab FTS-20 Fourier transform spectrometer. Results and Discussion The effect of an increasing level of solvation by ACN, THF, and MezSO on the value of Av3 for Li+N03-is apparent from Figures 1-3. Increasing the amount of solvent in the matrix increases the number of solvent molecules coordinated to the lithium cation. As was thoroughly demonstrated for NH3 and H20in earlier studies, it is clear that the addition of aprotic molecules of solvation also
The Journal of Physlcal Chemistry, Vol. 85, No. 10, 1981 1389
Li+N03- in Aprotic Solvents
TABLE I: Observed and Derived Spectroscopic Values for the Nitrate Ion of the Solvates of Li'N0,- a solvent Ar ACN THF Me SO
m
0 1268 1530 26 2 0
3 4 1 2 3 1 2 3 4 1327 1341 1287 1298 1325 1292 1303 1327 1335 1440 1400 1468 1430 1496 1492 1472 1453 1418 v3a 113 59 204 165 174 105 126 83 205 Av3 0.78 0.37 0.57 0.60 0.22 0.52 0.68 0.22 0.34 6v3 0.18 0.19 0.20 0.20 0.22 0.17 0.17 0.22 0.17 0 6v,lm Here v S 8 and V 3 b , given in units of cm-', refer t o the split components of the anion v3(e)mode and m is defined by the formula Li'NO, -.(solvent),. The 6 v 3 / mvalues are accurate to *0.01 units. ab
1550
1 1285 1503 218 0.17 0.17
2 1300 1473 173 0.34 0.17
1350
1450
125
cld
Flgure 3. Infrared bands for the v3 mode of Li+N03- isolated in matrices with compositions that vary from pure argon to pure Me2S0. The superimposed numerals identify v, doublets with the number of solvating Me2S0 molecules per ion pair. The Me,SO estimated matrix concentrations in mole percent, corresponding to curves from top to bottom, are 0.0, 1.0, 8, 50, and 100.
causes a discontinuous reduction in the degree of cation distortion of the anion and, consequently, a stepwise decrease in the value of Av3. The novel aspect of the new data derives, as anticipated, from the relatively sharp quality of the v3 band components for the nitrate ion. Because of this property the overlapping of bands for different levels of solvation is minor, particularly for V3b, so that the assignment of pairs of band frequencies to the two v3 components (v3,, V3b) of a particular solvate, Li+NO,-.(solvent),, becomes obvious from data such as presented in Figures 1-3. This assignment, which leads directly to firm values of the maximum solvation number n and which reveals the relative effect of each solvate molecule in a complete solvation sequence, is indicated by the labeling of prominent band features in Figures 1-3 and is summarized in Table I. There is a minimum of uncertainty in these assignments since bands assigned to a given solvate appeared and disappeared in predictable patterns as the
matrix concentration of a particular solvent was varied. The greatest uncertainty was met in the assignment for the first MezSO solvate (Li+NO3-.MeZSO)and was associated with the special difficulty in maintaining the anhydrous quality of the deuterated Me2S0. The bands labeled (1) in Figure 3 were initially assigned to the first hydrate of LiN03 (Li+NO3--H20),but the results for two complete and separate series of experiments employing different MezSO transfer and drying procedures are most consistent with the assignment offered in Figure 3 and Table I. Further, it is clear that the solvation pattern for MezSOwould be radically different than for ACN or THF, if the pair of bands labeled (2) were assigned to the first MezSO solvate. The relative effect of successive solvate molecules on the cation ability to distort the nitrate anion can be deduced from the variation in the average fractional decrease in the value of Av3 per solvate molecule for a series of solvates, Li+N03-.(solvent),. For this reason values for the fractional decrease in Au3, bug, defined as [AY: - Av3(solvate)]/Av: where A u t is the pure argon matrix value (262 cm-l), and for 6v3/m, the fractional decrease per solvating molecule, are also included in Table I. It should be noted that the range of possible bv3 values is from 0.0 (maximally distorted anion in the isolated contact ion pair) to 1.0 (the symmetric undistorted Dsh anion). In analyzing the errors associated with the values assigned to Au3, du3, and 6u3/m in Table I, the limiting factor is the accuracy of the measurement of the spacing between the two v3 band componenta of a given solvate, rather than the measurement of the absolute positions of the two bands. Thus, the measured values of Av3 for the m = 3 solvate of ACN, MezSO, and THF are 126,113, and 105 cm-', respectively. An examination of the superimposed spectra for these three cases reveals an obviously significant difference in the Au3 values as reflected in the tabulated values in Table I. Further, the results for multiple sampling suggest that Av3 can be determined reproducibly to better than f 2 cm-l, provided the measurement for a given spectrum is restricted to the Au3 value for the dominant solvate. Spectra for numerous samples not represented in the figures have been obtained to achieve the desired accuracy and to test this reproducibility. The 6u3/mvalues are all determined to within fO.O1 unib and the reliability for the larger m values ( m 2 2) is much better than suggested by such limits. Solvation Numbers. The most basic information in Table I for each solvent reduces to a value for the maximum solvation number n for the LiN03 contact ion pair. For ACN and MezSO n assumes a value of 4 while n is 3 for THF. These are firm values deduced from spectroscopic data that can be given no other reasonable interpretation. Further, these numbers clearly apply to the corresponding liquid solutions since it has been shown, for ACN and THF, that the contact ion pair solution spectra are essentially identical with the ion-pair pure-glassymatrix It is interesting that the THF liquid solution spectrum did show a significant concentration of
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The Journal of Physical Chemistry, Vol. 85, No. IO, 1981
what can now be clearly identified as the 2-solvate (i.e., v3, = 1470 cm-'), along with the dominant 3-solvate, and the matrix spectra for pure glassy Me2S0 shows a dominance of the Csolvate, but also a significant concentration of the 3-solvate. Since the experience with THF suggests that there is a possible shift in favor of the higher solvate in the matrix case vs. the liquid solution, the possiblity of the 3-solvate being as important as the 4-solvate for Li+NO3- in liquid Me2S0 must be considered. However, in so far as the ion-pair v3-band components are observable in liquid Me2S0 solution, for which the principal nitrate species is the nonassociated anion, the frequencies agree with the values for the 4-solvate (i.e., vg. is near 1400 cm-', not 1440 cm-')." It is informative to compare the values of n determined spectroscopically with values reported from earlier studies. Before doing so, however, it must be appreciated that our spectroscopic n values are for the solvation of contact ion pairs, a fact that has been shown repeatedly in earlier ~ t u d i e s . ~Further, both spectroscopic and quantum chemical results point to a bidentate association between Li+ and NO, regardless of the degree of solvation.1b8c*6 Consequently, n must be increased to n 1,or possibly n + 2, to give the proper coordination number for bare Li+ which is, therefore, 5 (or 6) for ACN and Me2S0 and 4 (or 5 ) for THF solutions. These numbers are large when compared to most suggested values. For example, in a theoretical study based on a charge-polarizability model, Kruus and Poppe found n = 2 for bare Li+ in a Me2S0 solution? Kebarle has found that, after the hydration number of the gaseous Li+ ion reaches 4, additional water molecules apparently do not enter the inner hydration sphere" while Clementi and Barsotti have also deduced a hydration number of 4 from a theoretical approach.la Thus, even though the nitrate ion in Li+N03- might be replaced by a single solvent molecule in the case of the dissociated ions, the most reasonable values for the maximum solvation number of bare Li+, as determined by an extrapolation from our spectroscopic results for contact Li+N03-, are somewhat greater than anticipated from other studies. Soluation Effects. A remarkable result that elucidates the strength, and less directly the nature, of the interaction of the mth solvating molecule with the Li+NO< ion pair is the constancy of the value of 6v/m for the solvent ACN (i.e., 0.17) regardless of the value of m (Table I). This result suggests that for this case, for which steric factors should be of minimum importance, considering the near linear structure of the ACN molecule, the effect of each successive ACN solvate molecule on the bonding in the NO3- anion is as significant as that of the first solvate molecule. This result is remarkable because the maximum value of bv3 is 1.0 (zero anion distortion) and, having achieved a value of 0.71 when m = 4, the effect must saturate rapidly, particularly should the imaginary case of m > 4 be realized. In fact, in the absence of saturation, 6v3 would be 0.88 for m = 5. This approaches closely the 6v3 = 1.0 undistorted nitrate-ion case and, therefore, implies that the cation effect (charge) on the anion would approach zero in the m = 5 solvate. Since the cation interaction with the solvent and the magnitude of the anion distortion depend on similar factors (e.g., cation charge density or polarizing power) it is clear that the attraction of the cation for additional solvating molecules
+
(5) J. A. Draeger, G. Ritzhaupt, and J. P. Devlin, Znorg. Chem., 18, 1809 (1979). (6) P. Kruus and B. E. Poppe, Sixth International Conference on Non-Aqueos Solutions, Waterloo, Ontario, Aug 1978, paper V-7.
loth et al.
must decrease sharply for m > 4. A corollary result is that, contrary to intuition, one can expect solvation numbers to tend to decrease with superior solvent donicity. Since 6v3/m will be greater the stronger the donor solvent molecule interaction with the cation, the near saturation of the range of possible 6v3 values (0-1.0) will occur more quickly for solvents having superior donor strength. This establishes a perspective on the tendency of both THF and Me2S0 to aasume values of n = 3, rather than n = 4, with Li+N03-. From Gutman's assignment of donicity numbers THF and Me2S0 rate well above ACN in donor ability? This is reflected in average 6v3/mvalues of -0.20, rather than 0.17, for both solvents. When m = 4 this gives bv3 0.80, a value which must correspond closely to the solvate instability point which occurs as the charge of the cation is reduced by solvation. Since, in fact, MezSO prefers m = 4 while THF does not (although the donicity number of M e a 0 is greater than for THF) other factors such as steric hindrance or solvent molecular polarity must also have a significant influence! Such factors will be stressed in a subsequent paper that will include spectroscopic data for stepwise solvation of Na+N03- and K+NOc by ACN, THF, Me2S0, and dimethylformamide. As noted above steric factors should be of minimum importance for ACN. However, THF and, to a lesser degree, Me2S0 are space filling about the oxygen coordination site of the solvent molecules so steric effects must be expected. In this respect the fluctuation in the value of (Sv3/m) with increasing m value for THF and Me2S0 (Table I), contrasted with the constancy of that value for ACN, may be significant. Despite these fluctuations, the magnitude of 6v3/mis little changed for the complete range of m values, again indicating that, as the solvation number is increased, the effect of each solvating molecule on the cation polarizing power remains large and comparable to the effect of the first molecule of solvation. In making comparisons between THF and Me2S0 the most distinguishing relevant experimental fact is that LiN03dissociates in liquid M e a 0 but not in a liquid THF solution." Since v3(e) for the NO, of the dissociated salt in Me2S0 is not split, it is unlikely that solvent-anion interactions are sufficiently great to be the cause of this difference. That being the case the data in Table I, which embody detailed information regarding the solvent-cation interaction, should be informative. The solvation effects of THF and Me2S0 molecules, represented by bv3/m, for corresponding m values are nearly identical for m = 1to m = 3. The notable difference for the two cases is that Me2S0 achieves the m = 4 state while THF does not. Why? Steric factors, namely the greater filling of space about the coordinating solvent oxygen atom in THF seems to be the most reasonable, but nevertheless tentative, answer. That is, the greater ability of Me2S0, relative to THF, to displace the anion from its contact position in liquid solution may merely reflect the greater ease of packing 4 or more MezSO molecules into the inner solvation sphere and may be unrelated to solvent polarity or donicity. Obviously this steric effect cannot explain the difference between Me2S0 and ACN (which, like THF, solvates Li+NO< without dissociation) so in the latter comparison the greater donicity of MezSO is identified as the critical factor. (7) V. Gutmann, "The Donor-Acceptor Approach to Molecular Interaction", Plenum Press, New York, 1978. (8) For a detailed analysis of the deficiencies of the single scale donicity number approach see R. s. Drago, Pure Appl. Chem., 52, 2261 (1980). (9) T. J. V. Findlay and M. C. R. Symons, J. Chem. SOC.,Faraday Trans. 1, 72, 820 (1976).
J. Phys. Chem. 1861, 85, 1391-1395
The values of 6v3 listed in Table I tempt one to conclude that the association or nonassociation of ion pairs in liquid solution is linked to the charge redistribution that is directly measured by the value of 8v3,1bwith the crossover value for Li+N03-lying somewhere between 0.71 and 0.78. Results reported for ammonia may suggest a more complicated picture. The value of 6v3 for fully solvated LiN03 a value that is 0.82,as measured in glassy NH3 is consistent with the Raman spectra which, for dilute solutions, are dominated by nonassociated nitrate ions.l0J1 However, conductance measurements12of LiN03 in liquid ammonia suggest that ion pairs and higher aggregates are the predominant species in this solvent, a result not expected from the value of 6v3 alone. Apparently bulk solvent effects, as well as molecular parameters, such as those reflected by 8v3, affect the degree of dissociationof ion pairs in liquid solution. For solvents of comparable dielectric constant, however (e.g., Me2S0, ACN), cation charge cancellation, as reflected in the value of 6v3, seems to be the determining factor. (IO) D. J. Gardiner, R. E. Hester, and W.E. L. Grossman, J. Chem. Phys., 69, 175 (1973). (11) P. Gans and J. B. Gill, Faraday Discuss. Chem. SOC., 64, 150 (1978). (12) J. J. Lagowski and G. A. Moczygemba in “The Chemistry of Non-Aqueous Solvents”, Vol. 11, J. J. Lagowski, Ed., Academic Press, London, 1967.
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Conclusions The results of this study can be summarized as follows: (a) The nearly linear ACN molecule, which is a relatively weak electron donor, uees four molecules to solvate the lithium cation in Li+N03-. The effect of each solvating molecule is constant and, reflecting the small ACN donicity number, relatively small ((6v8/m) = 0.17). (b) The stronger electron-donor solvent Me2S0 also has a maximum solvation number of 4 with Li+N03- but roughly a third of the Li+N03-ion pairs in pure glassy MezSOaccept only three solvate molecules. Consistent with a donicity number of -30, the effect of the first and subsequent MezSO solvating molecules is large ((6v3/m) = 0.20). (c) THF, which has a donicity number closer to ACN than MezSO,behaves more like MezSO as it solvates Li+N03-. In particular, the value of (6v3/m) is very similar for the two solvents for each m value. However, THF has a maximum solvation number of 3 and, in liquid THF, approximately one-third of the Li+N03-ion pairs accept only two solvent molecules. These results seem to reflect both donicity and steric effects as discussed in the Results and Discussion. Acknowledgment. This research was supported by the National Science Foundation under grants CHE 77-09653 and CHE 79-25567.
Laser Flash Photolysis Study of Electron Transfer Reactions of Phenolate Ions with Aromatlc Carbonyl Trlplets‘ P. K. Das’ and S. N. Bhattacharyya? Radiation Laboratory, University of Notre Dame, Notre Deme, lndlena 46556 (Received: December 2, W60)
Light-induced electron transfer reactions from phenolate ions to a variety of carbonyl triplets have been studied in aqueous acetonitrile by using 337.1-nm laser flash photolysis. In particular, the effects of carbonyl triplet nature and substitution in phenolate ions have been examined with respect to both kinetics and primary photoproduct yields. The rate constants for electron transfer are found to be in the range 2 X log-1 X 1Olo M-‘s-l. The quantum yields of primary photoproducts (phenoxy radicals and radical anions of carbonyl compounds) are essentially unity except with p-bromo- and p-iodophenolate ions. With p-iodophenolate ion, the yields of ketone-derivedradical ions are in the range 0.3-0.5 and remain practically unchanged on increasing the temperature from 23 to 75 “C.This result suggests the involvement of heavy atom induced intersystem crossing and/or back electron transfer at some stage(s) of the electron transfer reaction mechanism.
Introduction Light-induced redox reactions are becoming subjects of increasing importance these days. The current interest in these reactions stems primarily from the fact that they are potentially useful in the utilization of solar energy. A large number of photogenerated transients such as radicals, biradicals, radical ions, and singlet and triplet excited (1)The research described herein waa supported by the Office of Baaic Energy Sciences of the Department of Energy. This is Document No. NDRL-2190 from the Notre Dame Radiation Laboratory. The paper waa presented in part at the Symposium on Excited States in Chemistry and Biology held in March, 1981, at IIT, Delhi. (2) On leave of absence from Saha Institute of Nuclear Physics, Calcutta, India. 0022-3654/81/2085-1391$01.25/0
states have been studied and characterized for their redox behaviors. The precursors of these transients encompass a wide variety of systems3 ranging from inorganic complexes to organic dyes, carbonyl compounds, and aromatic hydrocarbons. (3) Some recent representative references are the following: (a) Ohno, T.; Lichtin, N. N. J. Am. Chem. SOC. 1980,102,4636-43. (b) Scheerer, R.; Griitzel, M. Ibid. 1977,99, 865-71. (c) Wilkinson, F.; Schroeder, J. J. Chem. SOC.,Faraday Trans.2 1979,441-50. (d) Hoeelton, M. A.; Lin, C.-T.;Schwarz, H. A,; Sutin, N. J.Am. Chem. SOC.1978,100,2382-8. (e) Kuzmin, V. A.; Darmanyan, A. P.; Levin, P. P. Chem. Phys. Lett. 1979, 63, 610-4. (0Vogelmann, E.; Schreiner, S.;Rauscher, W.; Kramer, H. E. A. 2.Phys. Chem. (Wiesbaden) 1976,101,321-36. (g) Scaiano, J. C.; Lissi, E. A.; Encinaa, M. V. Rev. Chem. Intermed. 1978,139-96. (h) Osif, T. L.; Lichtin, N. N.; Hoffman, M. 2.;Ray, S. J. Phys. Chem. 1980,84, 410-4.
0 1981 American Chemical Soclety
