10359
J. Phys. Chem. 1991, 95, 10359-10369 have to be tested for self-consistency in more realistic modelsultimately by experimental verifications. In accord with previous models of heavy-at~m-blocking~~"~ and sequential IVR,70,71the next step in this direction should extend the present investigation from simple inversion to coupled inversion-rotation of the FePH2 group, possibly including the PH stretches. In any case, the present results point to series of IR picosecond laser pulses with analytical shapes as novel, promising tool for laser-assisted chemistry.
Acknowledgment. We extend thanks to Professor W. Malisch for stimulating discussions about his [ C P ( C O ) ~ F ~ P complex, H~] to Mrs. E. Kolba for presenting our results at the Kasha conference "Photoinduced Proton Transfer Dynamics", Tallahassee, FL, Jan
1991, and to Professors P. Brumer, M. Shapiro, and M. Quack for sending their papers prior to publication. G.K.P. thanks Mr. T.-M. Kruel for emergency computational support during some midnight sessions. J.M. thanks Professor X. Chapuisat for indepth correspondence on related aspects of the multidimensional dynamics in the system, and President M. Gorbachev for catalyzing this international cooperation. Generous financial support by the Fonds der chemischen Industrie and by the Deutsche Forschungsgemeinschaft (project DFG-SFB 347/C3) is also gratefully acknowledged. The computations were carried out on the Cray Y M P at HLRZ, Julich, VAX 6000-410 and S U N at the Rechenzentrum der Universitat Wurzburg, and our SUN and micro-VAX-I1 computers.
7-Azaindole in Alcohols: Solvation Dynamics and Proton Transfer Richard S. Moog* Department of Chemistry, Franklin & Marshall College, Lancaster, Pennsylvania 17604-3003
and Mark Maroncelli* Department of Chemistry, 152 Davey Laboratory, The Pennsylvania State University, University Park, Pennsylvania 16802 (Received: March 15, 1991; In Final Form: July 1 , 1991)
The photoinduced excited-state double proton-transfer reaction of 7-azaindole has been examined in a variety of alcohol solvents. The influence of temperature and solvent deuteration has been investigated. A time-dependent Stokes shift of the initially excited normal species is observed, and this species is found to be the kinetic precursor to the tautomeric form. A substantial overlap of the normal and tautomer emissions is found, indicating that the tautomer emission must be monitored at wavelengths of 550 nm or greater to avoid contamination from the normal emission. The observed proton-transfer times in alcohols at room temperature are well correlated with the solvation parameter ET(30), suggesting that the rate of proton transfer is related to the strength of solventsolute interactions. However, an unusual temperature-dependent isotope effect is also observed, in which the relative rates of proton transfer in normal and deuterated alcohols become closer as the temperature is lowered. These results are interpreted in terms of a two-step model for proton transfer, involving solvent rearrangement to an appropriate configuration for the reaction to occur, followed by rapid proton transfer.
I. Introduction The role of the solvent in the excited-state dynamics of 7azaindole (7AI) has been a subject of some interest and controversy since the discovery of its photoinduced double proton-transfer reaction over twenty years ago. In the seminal work on this reaction, Kasha and co-workers' proposed that dimers of 7AI in alkane solvents and 7AI-alcohol complexes in alcohol solutions can undergo a photoinduced excited-state double proton transfer to produce a tautomeric species. (See Scheme I.) In such solutions, two emission bands are generally observed, as illustrated by the spectrum in methanol solution shown in Figure 1. The more intense, higher energy emission is due to species which have not undergone proton transfer and has been labeled the normal emission (N), while the lower energy emission peaked around 500 nm has been attributed to the tautomeric species (T).1-3 The rate of proton transfer in 7AI dimers has been shown to occur on the picosecond time ~ c a l eover ~ , ~a wide range of temperatures, and the relative amount of tautomer emission is essentially in~~~~
SCHEME I
:
I
I
R
I
R
~~
( 1 ) Taylor, C. A,; El-Bayoumi, M. A,; Kasha, M. Proc. Nutl. Acud. Sci. U.S.A. 1969, 63, 253. (2) Ingham, K. C.; Abu-Elgheit, M.; El-Bayoumi, M. A. J . Am. Chem. SOC.1971, 93, 5023. (3) Collins, S . T. J . Phys. Chem. 1983, 87, 3202. (4) Hetherington 111, W. M.; Micheels, R. M.; Eisenthal, K. B. Chem. Phys. Lett. 1979, 66. 230. (5) Share, P. E.; Sarisky, M. J.; Pereira, M. A,; Repinec, S. T.; Hochstrasser, R. M. J . Lumin. 1991, 48/49, 204.
dependent of temperature below 150 K.2,6 However, in alcohols the production of tautomers through this excited-state reaction is slower than in the dimer and has a stronger temperature dependence. The activation energy of tautomerization, determined from the relative intensities of the normal and tautomer emissions, ( 6 ) Ingham, K. C.; El-Bayoumi, M. A. J. Am. Chem. SOC.1974,96, 1674.
0022-3654/91/2095-10359%02.50/0 0 1991 American Chemical Society
Moog and Maroncelli
10360 The Journal of Physical Chemistry, Vol. 95, No. 25, 1991
has been observed to correlate well with the activation energy for viscous flow in alcohol^.^-'^ These results have been interpreted as suggesting that some large-amplitude solvent motion relative to the 7AI molecule is necessary for the proton-transfer reaction to occur in alcohols. Thus, solutions of 7AI in alcohols may provide an interesting system for the study of the interplay between solvent dynamics and chemical reaction dynamics. To better understand the dynamical processes occurring in 7AI alcohol solutions following photoexcitation, time-dependent fluorescence spectroscopy has proven to be a useful tool. The first such work on the picosecond time scale was reported by McMorrow and Aartsma'l in 1986, who observed single-exponential decay of the normal emission in methanol at 395 nm, but a two-component rise to the emission at 510 nm. These results were interpreted in terms of a ground-state equilibrium involving two distinct types of solvent arrangements around the 7AI molecules. They proposed that some 7AI molecules are cyclically hydrogen bonded with a single alcohol molecule as shown in Scheme I, enabling them to tautomerize directly and rapidly (26the decay of the normal species must be predominantly through nonradiative processes. The question thus becomes one of determining whether some nonradiative deactivation process other than proton transfer contributes significantly to the depopulation of the normal excited state. Electron ejection from an exciplex has been proposed previously as a possible candidate for this dea~tivation,~ based on steady-state fluorescence spectroscopy of 7AI in mixed alcohol/ alkane solvents. The time-resolved measurements reported here allow a further quantitative analysis of this question, which is presented below. For the tautomer, the overall deactivation rate constant is kT and the radiative rate constant is kTad. Thus, the quantum yield of fluorescence for the tautomer may be defined in the standard manner as 4T= kTad/kT. Note, however, that this is not in general obtainable by measuring the relative intensity of the tautomer emission in the steady-state spectrum, because the initially excited species there is not the tautomer. The quantity which may be calculated based on the relative intensity of the observed tautomer emission is @T', defined as the ratio of the number of photons emitted by the tautomer to the number absorbed by (nwmal) 7Al. Note that 4 ' = OT only if k = kpt. On the other hand, if kpt < k , then a significant number of originally excited 7AI molecules will not be converted to tautomers, and the result will be that 4T'
< 4T.
The quantity 4T'can be estimated from the steady-state fluorescence spectrum using standard techniques for measurement of fluorescence quantum yield^.'^^^' Although the normal and tautomer bands overlap substantially, a reasonable estimate of the tautomer emission intensity may be obtained from the integrated intensity for wavelengths above the spectral minimum, which appears at about 450 nm. (Considering the relatively large uncertainties in the value of @T as described below, a more precise analysis of the intensity is not justified.) These values are shown in Table 1 for several representative alcohols. The value for @T may be calculated from the values of kT and kTad. The total deactivation rate constant is simply 1 / ~ ? ,obtained from the time-resolved measurements, and is listed in Table I. The value for kFad is more difficult to ascertain. The best estimate is provided (25) El-Bayoumi, M . A.; Avouris, P.; Ware, W. R. J . Chem. Phys. 1975, 62, 2499. ( 2 6 ) Avouris, P.; Yang, L. L.; El-Bayoumi, M . A. Photochem. Photobiol. 1976, 24, 211.
(27) Ediger, M . D.;Moog, R. S.;Boxer, S.G.; Fayer, M . D. Chem. Phys. Letr. 1982, 88, 123.
Moog and Maroncelli
10364 The Journal of Physical Chemistry, Vol. 95, No. 25, 1991 TABLE I: Comparison of Tautomer Quantum Yields Observed and Calculated Assuming k = k, (293 K)
solvent methanol methanol-0-d ethanol 1-decanol 2-methyl-2-propanol 2,2,2-trifluoroethanoI ethylene glycol
kT.a ns-l
1.51 1.20 1.33 0.76 0.96 2.57 1.37
102&0lb 1.2
1.8 1.2 2.1 1.2
2.4 3.5
3.1 5.5
0.2
3.0 1.5
1.4
3.3
r--
400’
103dT‘c 103bTd
3
f-
2.6 3.3 3.0 5.3 4.2
“Total decay rate of the tautomer, kT = Tdw-l (550 nm). bTotal fluorescence quantum yield determined relative to quinine sulfate in 1.0 N H2S04(4 = 0.546). Demas, J. N.; Crosby, G. A. J. Phys. Chem. 1971, 75,991. CMeasuredquantum yield of the tautomer under the assumption that all N* roduce tautomer. The value is determined dX]@t. dTrue fluorescence quantum from q5T’ {.f:$(X) dX/.f,&(X) ,El yield of the tautomer #T = kTrad/kT. For kTradwe have used the value 4 X lo6 s-l reported for the tautomer analogue 7-methyl-7if-pyrrolo[2,34]pyridine by: Waluk, J.; Pakula, B.; Komorowski, S.J. J. Photochem. 1987, 39, 49. Note that the values listed in this column should be considered to have only one significant figure due to the uncertainty in this value of kTrad. by the radiative rate of the analogous methylated species, 7methyl-7H-pyrrolo[2,3-b]pyridine (7MPP). This value has been reported by Waluk and co-workersZ8to be 4 X lo6 s-l. Table I shows the obtained values of 4Tand 4T’in five alcohol solvents. The agreement between these values is excellent, especially considering the uncertainty in the value of k;rad. Thus, the protontransfer reaction appears to provide the main deactivation pathway for the normal species in alcohol solutions, and is the predominant process whose rate is measured by the rise of tautomer fluoresence. Therefore, the rise time of tautomer fluorescence is taken to be an excellent approximate measure of the overall rate of the proton-transfer reaction: T , i= Ilkp, = 7pt. C. Solvent Dependence at Room Temperature. The rate of the proton transfer reaction depends significantly on the particular alcohol used as solvent. Table I1 shows the observed tautomer rise times, T ~ , at , 550 nm for a variety of alcohols at room temperature, along with several solvent parameters. The reported rise times are representative of the total deactivation rate of the normal species, and as discussed above, these values are taken as measures of the overall rate of the proton-transfer reaction.29 As can be seen from the data in Table 11, T~~ does not correlate in any simple way with measurs of solvent dynamics such as viscosity or the longitudinal relaxation time, 7L. Note that there is only a 25% increase in T~~ from 1-propanol to 1-decanol, while the viscosity increases about the 6-fold. On the other hand, the viscosity of 2-propanol is approximately the same as 1-propanol, but the observed 7ptis almost as large as that of 1-decanol. This absence of a systematic relationship between the rate of proton transfer and viscosity is also shown graphically in Figure 6a. However, with the exception of trifluoroethanol, this plot does show a general trend of increasing 7pl with increasing viscosity within a given class of alcohols. Measures of equilibrium solvent properties, such as the hydrogen bond donating ability a (Table 11) and “polarity” ET(30)(Table I1 and Figure 7a) are better able to correlate 7p, in the different alcohols. With the exceptions of the polyalcohols (ethylene glycol and glycerol), there is a good correlation between 7pt and a,and an excellent one with ET(30). Previous workers have similarly noted that the observed rate of the proton-transfer reaction is related to the tendency of the alcohol to donate its hydroxylic p r ~ t o n however, ;~ the nearly linear correlation with ET(30)for the entire range of monoprotic alcohols is surprising. The ET(30) parameter seems to provide a very accurate measure of the strength of the 7AI-alcohol interactions important for reaction. (28) Waluk, J.; Pakula, B.; Komorowski, S . J. J . Photochem. 1987, 39,49. (29) The minor correction to this value for the radiative rate is ignored because it is quite small and within the experimental uncertainty in the reported values. The corrections for the nonradiative deactivation rate of the
normal species are not made because they are also expected to be small, and they are difficult to ascertain.
1
1500.
1.6 2.9
1
e’ h 2
~~
-\
750.
3
1
0. 0.0
7.0
21 .o
14.0
V i s c o s i t y ( cP)
Figure 6. (a) Proton-transfer times and (b) tautomer lifetimes (ps) of 7AI in various alcohols (293 K) plotted as functions of solvent viscosity.
The numbering scheme differentiates between primary alcohols (1, l’), secondary and tertiary alcohols (2); and polyalcohols (3) as detailed in Table 11.
1500. 1 7
2
1
3
1
1
1’
0.’ 43.
’
’ 47.
’
’ 51.
’
’ 55.
’
’ 59.
’
~
63.
Figure 7. (a) Proton-transfer times and (b) tautomer lifetimes (ps) of 7AI in various alcohols (293 K) plotted as functions of solvent “polarity”, E~(30)(see text). The numbering scheme differentiates between primary alcohols (1, l’), secondary and tertiary alcohols (2); and polyalcohols (3) as detailed in Table 11. The datum for 7pf of glycerol is off the scale of this plot (ipt 600 ps) but is shown as 3! here for the sake of com-
pleteness.
-
Since in alcohols the ET(30) “polarity” mainly reflects the hydrogen bonding ability of the solvent, it is reasonable to conclude that the proton-transfer reaction rate varies with alcohol according to the strength of the hydrogen bond formed between 7AI and the solvent. D. Temperature Dependence and Isotope Effects. In order to learn more about the mechanism of the proton-transfer reaction in 7A1, the kinetics in several deuterated solvents have been examined. Table I11 compares the overall proton-transfer times ( 7 J and tautomer decay times (711)observed in five deuterated alcohols with their perprotio counterparts at 293 K. For the pairs of normal and deuterated solvents studied, the ratios of protontransfer times, p,,(D/H) = 7pt(D)/7pt(H).are in the range 2-4 at room temperature. These values are relatively modest for a primary isotope effect,30 especially when one considers that both the indolic H and the solvent are deuterated in ROD solvent^.^' (30) See for example the review by: O’Reffall, R. A. M. In Proton Transfer Reacrions; Caldin, E., Gold, V., Eds.; Chapman and Hall: London, 1975; p 201.
The Journal of Physical Chemistry, Vol. 95, No. 25, 1991 10365
7-Azaindole in Alcohols
TABLE 11: Solvent Properties, Proton-TransferTimes, and Tautomer Lifetimes of 7AI in Various Alcohols (293 K) solvent methanol ethanol 1-propanol 1-butanol 1-pentanol 1-octanol 1-decanol 2,2,2-trifluoroethanoI 2-propanol 2-butanol 2-methyl- 1-propanol 2-methyl-2-propanol ethylene glycol glycerol
class 1 1 1 1 1 1 1 1' 2 2 2 2 3 3
n," CP
ffb
0.59 1.17 2.32 2.94 3.68 8.95 13.5 2.00 2.43 3.63 3.91 5.94 19.9 1410
0.98 0.86 0.80 0.79
1.51 0.78 0.62 0.92
Ed30)' 55.5 51.9 50.7 50.2 49.1 48.3 47.6 59.5 48.6 47.1 49.0 43.9 56.3 57.0
71
-
? PS
PS 124 169 198 213 204 224 248 30 234 262 236 351 339 -600
7n.C
7,1ge
9.2 35 87 119 190
-
540 (25 "C) 93 125 91
-
80 -110
PS
664 754 882 966 1057 1225 1315 389 902 1010 993 1044 729
oSolvent viscosity, from: Riddick, J. A,; Bunger, W. B.; Sankano, T. K. Organic Solvents; Wiley: New York, 1986. bSolvent hydrogen bond donor acidity parameter (associated with the solvatochromic A* polarity scale of Taft and co-workers; from: Kamlet, M. J.; Abboud, J. L. M.; Taft, R. W.; Prog. Phys. Org. Chem. 1981, 13, 485. CSolvent "polarity" as measured by the solvatochromic scale of Reichardt and co-workers; from: Reichardt, C. In Molecular Interactions; Ratajczak, H., Orville-Thomas, W. J., Eds.; Wiley: New York, 1982; Vol. 3, p 241; and Reichardt, C.; Eschner, M.; Schafer, G.; Leibigs Ann. Chem. 1990, 57. dSolvent longitudinal relaxation times, determined from dielectric parameters via T~ = (c,/c,,)T~. In cases where more than one dielectric dispersion regime is observed, the 7L listed refers to the main (lowest) frequency regime. Values for methanol through 1-butanol and 2-propanol were obtained from the compilations of Castner, E. W., Jr.; Bagchi, B.; Maroncelli, M.; Webb, S. P.; Ruggiero, A. J.; Fleming, G. R. Ber. Bunsenges. Phys. Chem. 1988, 92, 363. The remaining data are from the following sources: 1-pentanol and 1-decanol: Dutuit, Y.; Salefran, J. L.; Bottreau, A. M.; Chahine, R.; Bose, T. K. Adv. Mol. Relax. Interact. Processes 1982, 23, 75. 2-Butanol and 2-methyl-1-propanol: Danhauser, W.; Cole, R. H. J . Chem. Phys. 1955, 23, 1762. Ethylene Glycol: Salefran, J. L.; Marjat, CI.; Vicq, G. Adu. Mol. Relax. Interact. Processes 1981, 19, 97. Glycerol McDuffie, G. E.; Litovitz, T. A. J . Chem. Phys. 1962, 37, 1699. eProton-transfer times and tautomer fluorescence lifetimes were taken to be the rise and decay times measured at 550 nm. These numbers are averages of 2-3 measurements made on separate days. The standard deviation in repeated measurements in *15 ps.
TABLE III: Comparison of Proton-Transfer Times and Tautomer Lifetimes in Several ROH/ROD Solvent Pairs (293 K)
MeOH/MeOD EtOH/EtOD 2-PrOH/2-PrODc l-BuOH/ 1-BuOD (CH~OH)~/(CHZOD)~
124 173 234 213 333
368 47 1 (990) 520 69 1
3.0 2.7 (4) 2.4 2.1
663 77 1 902 966 749
883 1021 (1300) 1264 1040
1.3 1.3 (1.4) 1.3 1.4
"The solvents are MeOH = methanol, EtOH = ethanol, 2-PrOH = 2-propano1, 1-BuOH = 1-butanol, and (CH20H)z = ethylene glycol. Deuterated solvents were substituted only at the hydroxyl positions. bProton-transfer times and tautomer lifetimes (ps) were taken to be the rise and decay times measured at 550 nm. The numbers reported are averages of 2-3 independent measurements. The standard deviation in repeated measurements was f 1 5 ps. The p(D/H) are the ratios T ( D ) / ~ ( H ) . 'The deuterated sample of 2-propanol contained significant amounts of a fluorescent impurity and these values should be regarded as subject to larger uncertainties than the rest.
*- BuOH
7 3.0 1
6.5 /7
'-(CH,
OH),
W
!-
I-
\ n
W
0
9 W -\
t
2.9
3.1
3.3
3.5
3.7
3.9
1000/T
2.8
3.2
3.6
4.0
4.4
1000/T Figure 9. Isotopic time constant ratios, 7(D)/7(H), in several solvents
Figure 8. Arrhenius plots of proton transfer ( T ~ , ,solid points and curves) and tautomer decay times ( 7 [ , , open points and dashed curves) of 7AI in EtOH (circles) and EtOD (squares). All times are in picoseconds. The lines shown are the linear fits to the data and correspond to the Arrhenius parameters provided in Table IV. The short-dashed curve also shown with the EtOD data is a quadratic fit to this data and is included merely to highlight the deviation from linearity.
as functions of temperature. The filled symbols refer to ratios of proton-transfer times and the corresponding empty symbols refer to the tautomer lifetimes in the solvents indicated. The error bars are indicative of the uncertainties in the 7pt ratios; the uncertainties in the ratios of T ~ are smaller than the symbols. BuOH refers to I-butanol.
However, such values are in line with isotope effects for other ultrafast proton-transfer reactions such a s t h e excited-state pro-
ton-transfer reaction of 3 - h y d r o x y f l a ~ o n e . W ~ ~h a t is most unusual about t h e proton-transfer kinetics of 7 A I in alcohols is t h e tem-
,
Moog and Maroncelli
10366 The Journal of Physical Chemistry, Vol. 95, No. 25, I991 7.1
'
'
,
'
'
'
'
'
'
'
'
'
'
'
'
SCHEME I l l
I
I H\O'..H
n
t-
u
I
6. "mcorrect' solvauon
- 5.
N
7. 6.
5. 2.8
3.2
3.6
4.0
4.4
1000/T Figure 10. Model fits to the experimental proton-transfer times. The points are the experimental proton-transfer times in ROH (circles) and ROD (squares) solvents. The model and parameters used to obtain the fits (solid curves) are described in the text. TABLE I V Arrhenius Parameters"Derived from the Proton-Transfer Times and Tautomer Lifetimes of 7AI in Several ROH and ROD Solvents T E,,! A E An, En', solvent range, K kJ/mol 101Bt61kJ/mol 1OIo sC1 kJ/mol
MeOH 213-333 MeOD EtOH 233-333 EtODe (CH,OH)2 283-343 (CH2OD)2
10.9 14.8 27.9
40 4.2 97 (3.7) 310 30
9.7 6.6 12.7 (7.1) 17.0 13.0
1.1 1.4 1.8 1.8 4.5 5.7
4.7 5.9 6.3 7.1 8.6 9.9
"Fits were made to the linear form In ( 7 ) = a + b / T (see Figures 8-10) and the parameters converted to A and E in the Arrhenius expression 7-l = A exp(-E/RT). Uncertainties in the A parameters are f 5 % . With the exception of the EtOD T~~ data, which is clearly nonArrhenius, the fits yielded correlation coefficients ( r 2 ) of better than 0.99 in all but one instance. bActivationenergies associated with solvent viscosity obtained by fitting data compiled in: Lundolt-Bornstein Physikalisch-Chemische Tabellen II; Springer: Berlin, 1969. Over the temperature ranges considered the solvent viscosities are well represented by an Arrhenius temperature dependence. Uncertainties in these coefficients of the linear regression are f0.3 kJ/mol for all but the EtOD data (which is clearly non-Arrhenius). duncertainties in these coefficients of the linear regression are less than f0.2 kJ/mol except for (CH20D2)2for which the uncertainty was f 0 . 6 kJ. 'The +pt data in EtOD is not well represented by an Arrhenius equation; see Figure 8. perature dependence of the deuterium isotope effect. Figures 8-1 0 display temperature-dependent data in normal and deuterated methanol, ethanol, and ethylene glycol. Protontransfer times are shown as filled symbols and tautomer lifetimes by open symbols; circles refer to ROH and squares to ROD solvents. With the notable exception of the temperature dependence of 7p, in EtOD, all data are reasonably represented by an Arrhenius equation over the temperature ranges studied. Activation energies and prefactors obtained from Arrhenius fits are provided in Table IV, as are activation energies associated with the solvent viscosity. Several features of these data are noteworthy. The activation energies for the proton-transfer reaction in the normal alcohols MeOH and EtOH are quite close to the viscosity activation energies of these solvents. While this correspondence does not hold in ethylene glycol, other workers have noted it in (31) NMR experiments demonstrate that in ROD solvents the indolic H exchanges upon mixing and so it is not possible to study the 7AI/D and RO-D isotope effects separately. Note that if the reaction were rate limited by a concerted two-proton exchange, then one would anticipate a squared isotope effect in deuterated alcohols that was in the range of 10-100 at room temperature. (32) Strandjord, A. J. G.;Barbara, P. F. Chem. Phys. Lett. 1983, 98,21.
'corecf 'solvauon
taufomer
N*
T
the former alcohol^^-^^ as well as one other normal alcohol, nb u t a n 0 1 . ~It~ ~is ~unlikely ~~ that this parallel between the temperature dependence of solvent viscosity and overall proton-transfer rate is merely fortuitous. These results imply the importance of large-amplitude motions in controlling the proton-transfer reaction. However, it is clear that such motions cannot be entirely rate determining, for if they were, no significant deuterium isotope effect would be expected, a prediction contrary to what is observed. Further, the variation of proton-transfer reaction rate with solvent at 293 K has been shown to be more a function of hydrogen bond donating ability of the solvent than of solvent viscosity. Therefore, the energetics of the actual proton-transfer step itself must also partly determine the reaction rate. The key to how solvent motions are coupled to the overall proton-transfer reaction is provided by the difference between the temperature dependence of the reaction rates in hydrogenated versus deuterated solvents. Whereas the reaction rate in EtOH (Figure 8) shows a nearly Arrhenius temperature dependence, the EtOD data exhibits an obvious curvature. In proton-transfer reactions, curvature in Arrhenius plots is often associated with the transition between a barrier crossing and a tunneling mechanism as temperature is lowered.33 However, in the present case, the curvature of the plot is in the direction opposite to that expected for such a transition. The apparent activation energy increases with decreasing temperature, contrary to the usual situation. An even more unusual feature is that the magnitude of the isotope effect decreases as the temperature decreases. The effect is directly illustrated in Figure 9 where the ratio T(D)/T(H) is plotted as a function of temperature. Such behavior is highly unusual and cannot be explained by a single-step proton-transfer mechanism. One other example where such behavior has been previously reported is in the proton abstraction from certain carbanions by methanol.34 In this case the kinetics were explained in terms of a multistep mechanism in which no single step was completely rate determining. A similar mechanism appears to be operative in the 7AIalcohol reaction. The unusual temperature dependence of the overall proton-transfer kinetics can be modeled using a simple two-step mechanism which is described in Section IIIE below. E. Two-step Model for the Overall Proton-Transfer Reaction. As originally proposed by McMorrow and Aartsma,' the overall proton-transfer reaction in 7Al-alcohol systems can be viewed as consisting of two steps: a solvent reorganization process and the actual proton-transfer event. A simple model for this reaction may thus be constructed. Based on the very fast proton transfer observed in 7AI dimer^,^ it may be assumed that if the hydrogen bonding arrangement is favorable, proton transfer is rapid. However, in contrast to the dimer case, 7AI molecules will exist in a wide range of hydrogen-bonded configurations in alcohol solvents; only a select subset of these configurations will be appropriately hydrogen bonded for the proton transfer to occur. Except for this 'korrect" subset of solvation states for which reaction is fast, proton transfer is assumed to be negligibly slow. The overall reaction kinetics observed for an ensemble of molecules will thus reflect both the rate of the intrinsic proton-transfer step for "correct" configurations and the rate of solvent reorganization by which "correct" arrangements are produced. To model this
'
(33) Kwart, H. Acc. Chem. Res. 1982, 15, 401. (34) Koch,H.F.; Koch, S. A. J. Am. Chem. SOC.1984, 106,4536. See also, Koch, H. F.; Dahlberg, D. B. J . Am. Chem. SOC. 1980, 102, 6102.
The Journal of Physical Chemistry, Vol, 95, No. 25, 1991 10367
7-Azaindole in Alcohols dynamics in a semiquantitative way we consider the approximate 2-step mechanism illustrated in Scheme 111. In this scheme, a single-state N is used to represent (excited) 7AI molecules in all solvation states inappropriate for proton transfer, N * represents those molecules in correct configurations, and T represents the tautomer. (The * = excited-state notation has been dropped here for simplicity.) In addition, k2 is the rate constant for the actual proton-transfer step; the rate constants k l and k-l reflect the dynamics of solvent reorganization. The kinetics derived from such a model can be obtained analytically using standard techn i q u e ~ . ) ~Assuming negligible initial population of correct solvation states (i.e., [N*], = o), the average rate of tautomer production36is given by
pair, the known solvent viscosity activation energy was used for El and the latter parameters were varied in order to find the best agreement between the temperature dependence of l/kavand the proton-transfer times. Since the effect of Kq is not actually independent of kl and k2,Kq was arbitrarily set at a value of 0.1. This particular value was used in light of the possible presence of =lo% prompt tautomer fluorescence in methanol. A brief discussion of the nature of the temperature dependence of this model is in order. At sufficiently high temperatures, the solvent reorganization step becomes fast enough so that k2 is rate limiting. In this regime the observed rates, k,,, for ROH and ROD solvents are
(r+-l - r--l)
kav =
(9) (4)
(r+-2 - r--2)
with 1 r* = -{(kl + k-l + k2) f [ ( k , + k-l + k2)2- 4klk2]1/2) (5) 2 This model may be used to explain the observed temperature dependence of the overall proton transfer reaction times rPl,which are equated with l/kav. Implementation of the model requires knowledge of the rate constants kl, k-,, and k2 as well as their temperature dependences, which are assumed to exhibit Arrhenius behavior. Thus, kl(T) = AI exp(-EI/kT) K,(T)
= kl(T)/k-l(T)
(6) (7)
= A , exp(-hEeq/kT)
M T ) = A2 exp(-WkT)
(8)
Six constants (Al, A,, A2,E l , AE? and E,) are therefore needed to completely specify the model. Since such detailed information is not available, some fairly severe assumptions must be made, with the goal of qualitatively explaining the unusual temperature-dependent isotope effect, rather than extracting any meaningful rate parameters from the model. The main assumptions are the following: (i) In analogy with the dimer behavior? the temperature dependence of k2 is likely to be weak and therefore E2is considered negligible in comparison to El. (ii) The activation energy for the solvent reorganization is set equal to the viscosity activation energy of the solvent, E l = E,. For methanol and ethanol such a choice is suggested by the closeness of the overall proton-transfer activation energy to E,,. (iii) The "equilibrium" activation energy, AE,, is assumed to be small compared to E l and thus Eq is approximately temperature independent. Although this may not be a particularly good assumption, it is used to reduce the number of unknown parameters. (iv) Finally, it is assumed that the only isotope-dependent rate constant is k2 and k2(D)/ k2(H) is independent of temperature (as it would be for a tunneling-dominated proton transfer). Given these four assumptions, the model can be completely specified in terms of the parameters k , , k2(H), k2(D), and Keq at one temperature, chosen to be 293 K. For each OH/OD solvent (35) See for example: Steinfeld, J. I.; Francisco, J. S.;Hase, W. L. Chemical Kinetics and Reaction Dynamics; Prentice-Hall: Englewood Cliffs, NJ, 1989. (36) The average rate is defined by
where [T(r)] is the concentration of excitated-state tautomer at time r assuming no loss of excited-state population. A few comments are in order regarding the use of this average rate constant for our purposes. Although the response is biexponential in the general case ([T(r)] = A+e-'+' A_e-N) for the parameter values of interest here, where k , , k-I> r- and A- is much greater than A+. From an experimental point of view, the response is indistinguishable from a singlaexponential response with rater-. Since k,, is within a few percent of r-, either rate could be used for characterizing the time dependence of the model. The average rate constant, k,,, is chosen to be consistent throughout all of the parameter space.
+
and the ratio k,,(ROD)/k,,(ROH) is temperature independent. In the opposite, low-temperature regime, the solvent dynamics dominate and the observed rate is independent of isotopic substitution, being simply k,( T). In the intermediate-temperature region the overall rates may show a non-Arrhenius temperature dependence and the ratio T ~ ~ ( D ) / T ~decreases ~ ( H ) with decreasing temperature in the manner observed experimentally. The 7AIalcohol reaction near room temperature is in this intermediate regime, in which neither step completely controls the reaction, and this fact gives rise to its interesting kinetics. Figure 10 shows the best fits achieved to the methanol, ethanol, and ethylene glycol data. The fits are not quantitative; however they do capture the main features of the experimental data. The rate constants used to produce these fits are (293 K): k l = 9, 6.5, and 4 ns-l; k2(H) = -500, -500, and -125 n d ; and k,(D) = 35, 35, and -30 ns-l for methanol, ethanol, and ethylene glycol, respectively. While the assumptions made are too crude to allow quantitative significance to be attached to these rate constants, the result that they all lie within reasonable ranges provides some justification for the applicability of the model. For example, the best fit yields proton transfer times of k2-l -2 ps in methanol and ethanol and -8 ps in ethylene glycol. These are all similar to the 1.4-ps time reported for 7AI dimer^.^ The isotope effect for the proton-transfer step k2(H)/k2(D) is 15 for methanol and ethanol and -4 for ethylene glycol, also within the expected range.37 Finally, the rates for the solvent reorganization step ( k , ) are ordered according to the relative solvent viscosities, as expected. Better agreement between the model and experiment, especially in the case of ethylene glycol, can be achieved by relaxing either the assumption El = E,, or by choosing AE, C 0. While either of these modifications could be justified, there is little basis for assigning other values to these parameters, and at present the simplest description seems the most appropriate. The model embodied in Scheme I11 thus provides a useful framework for interpreting the observed kinetics. Although this three-state description undoubtedly oversimplifies the problem, several conclusions should survive further refinements. First, the unusual temperature dependence of the overall proton-transfer reaction arises from the similar time scales of the solvent reorganization and the intrinsic proton-transfer steps such that neither completely dominates the overall rate. This idea is the main conclusion provided by application of the model. The observed behavior can only be explained if both the intrinsic proton-transfer step and the solvent reorganization step compete for control of the reaction rate. Second, in the protic solvents studied in detail, the reaction time observed is nearly equal to the solvent reorganization time. For example, for ethanol at 293 K, the solvent reorganization time (k,-I) is calculated to be 154 ps and the overall reaction time (kaY1)is 176 ps. This conclusion that the time scale of reaction in 7AI/ROH systems is essentially that of the solvation dynamics should be considered more robust than the simple model used to describe the kinetics. The fact that the activation energies
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(37) Based on ratios of the steady-state fluorescence of the N and T bands relative rates of H versus D dimers of between 10 (184 K) and 17 (132 K) were measured by: Tokumura, K.; Watanabe, Y.; Itoh, M.; J . Phys. Chem. 1986, 90,2362.
10368 The Journal of Physical Chemistry, Vol. 9.5, No. 25, 1991
for the overall proton-transfer reaction in methanol, ethanol, and l-butano17-10are very close to the respective viscosity activation energies implies that the rate of the intrinsic proton-transfer step must be much faster than the observed reaction rate (and relatively constant) over the temperature ranges considered. The same may not be true of all of the alcohols considered in section IIIA; it will be of interest to investigate the temperature dependence of a slower monoalcohol such as tert-butyl alcohol. One might question the interpretation of the observed reaction rates being essentially the rate of solvent reorganization on the basis of the lack of a linear dependence of 7pl on viscosity and the remarkably good correlation observed with the equilibrium solvation parameter ET(30). The nonlinear dependence of overall proton-transfer time with viscosity is not surprising. As mentioned previously, with the exception of tert-butyl alcohol and TFE, which have extreme values of ET(30),T~~ for the remaining monoalcohols does correlate reasonably well with viscosity (Figure 6a). Ignoring other changes, one could interpret Figure 6a as indicating that k , in this model is related to viscosity but in a nonlinear manner such that k , becomes “saturated” at higher viscosities. There are many examples of this type of behavior, where the microscopic friction that determines kinetics of small molecule solutes increases more slowly with chain length than does the bulk viscosity.38 Why the observed rates correlate so well with ET(30) is somewhat more puzzling. As previously discussed, ET(30) values may be regarded as a secondary measure of the strength of hydrogen bonding between 7AI and the various alcohols. A number of solvent properties related to this parameter will influence the rate of the overall proton-transfer reaction. First, the local hydrogen-bonding energetics measured by ET(30) may provide a better relative gauge of the time scales of local solventsolute reorganizations than does the bulk viscosity. That is, k , could be simply correlated to ET(30). (The fact that the polyalcohols deviate markedly from the correlation observed for the monoalcohols could be viewed as the breakdown of this relation between k , and ET(30) when the solvent is capable of hydrogen bonding at more than one site.) A related effect is that the nature of the distribution of solvent configurations surrounding the 7AI solute would be expected to be correlated with ET(30) and contribute to the solvent dependence. An increase in probe-solvent hydrogen bonding would imply a narrower distribution of solvent-solute configurations. If this narrowed distribution of populated solvation configurations (N) is closer to the correct (N*) configuration for proton transfer, then one would expect k, to be correspondingly increased. Although beyond the scope of the simple two-step model, such effects would arise naturally in the context of more sophisticated reaction models, such as those involving continuous diffusion in a reactive potential well.39 Finally, the model calculations suggest that even in the regime k2 >> k , , which is likely to apply to most of the protic alcohols, the intrinsic proton-transfer rate k2 still influences the reaction rate. For fixed k , and k-.,,variation of k2 over the range k2 = 10kl m causes the observed reaction time to vary by a factor of 2. Such changes could also be the source of the dependence of reaction time on ET(30) displayed in Figure 7a. Excluding trifluoroethanol and tert-butyl alcohol, 7p, for all of the monoalcohols studied varies by no more than a factor of 2 at room temperature and thus could be accounted for on the basis of variations of k2 alone. It seems reasonable that k2 should decrease with decreasing acidity or hydrogen bond donating ability of the alcohol involved and thus be correlated to ET(30) in the manner observed. F. Tautomer Lifetimes. The tautomer produced from the excited-state proton-transfer reaction exhibits fluorescence life, show interesting trends with solvent and temtimes, T ~ which perature. The relevant data for 7n are listed in parallel with the T ~ results , in Tables 11-IV and are illustrated in Figures 6-8. Based on previous studies of several alkylated tautomer ana-
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(38) See for example: Kim, S. K.; Fleming, G . R.; J . Phys. Chem. 1988,
92, 2 168.
(39) See the review by: Bagchi, B.; Fleming, G. R.; J . Phys. Chem. 1990, 94, 9.
Moog and Maroncelli logues,28the depopulation of the tautomer excited state is expected to be controlled mainly by internal conversion, not intersystem crossing to the triplet state. Thus, variations in 7n observed here reflect variations in the efficiency of internal conversion. The dependence of 7,, on solvent at room temperature can be seen in Table I1 and Figures 6b and 7b. In contrast to the behavior of T ~ there , is a much better correlation of q with solvent viscosity than with ET(30). This correlation with viscosity suggests that internal conversion is related to some sort of large-amplitude motion which modulates solute-solvent interactions. The lifetimes show a smaller isotope and temperature sensitivity than do the proton-transfer reactions rates. As shown in Table 111 the ratios pfl 7fl(D)/rfl(H)are all in the range 1.2-1.4 at room temperature. Since these values are close to 2Il2, it is tempting to ascribe the isotope effect to a difference in frequency factor associated with some high-frequency R-H stretching coordinate. However, pfl shows a significant temperature dependence, and, upon Arrhenius analysis (Table IV), the prefactors for H and D are found to be in the opposite order (Afl(H) < Afl(D)) from that expected. There is a substantial activation energy associated with the tautomer lifetime; it is in the range 1/3-1/2 of the viscosity activation energy of these solvents, or 5-10 kJ/mol. Finally, the difference between the activation energies for H and D isotopes is roughly constant (1.1 kJ/mol) among the three solvents studied. Unfortunately, there is no suitable theory with which to interpret the above results; however, a few observations may be made relating 7Al-alcohol systems to previously studied cases. The lifelimes of the 7AI tautomer in alcohols are not very different from those of tautomer analogues of the form 7RPP (7-R-7Hpyrrolo[2,3-b]pyridine) studied by Waluk et a1.28 These authors determined nonradiative decay times for three such compounds in which R = -CH3, -CH2COOCH3, and -CH2COC6H5. In methanol/ethanol mixtures (1 :4) at 293 K, decay times of 2,0.6, and 1 ns, respectively, were observed. These values are all near to the results shown in Table I11 for the 7AI tautomer in methanol and ethanol (0.6-1.0 ns). Moreover, Waluk et a1.28measured activation energies associated with these decay times to be 7.3, 6.6, and 7.8 W/mol, which are also similar to the values observed for the 7AI tautomer.40 In addition, Avouris et a1.26reported that the fluorescence quantum yield of the methylated analogue, 7MPP, was enhanced by a factor of 1.4 in ethanol-0-d relative to ethanol. All of these parallels between the 7AI tautomer and N-substituted analogues suggest that the presence of the N-H bond is not essential to the nonradiative decay mechanism. The present results may also be compared to those previously obtained for 7AI dimers in nonpolar solvents. Bulska et al.IO determined the lifetimes of the 7AI(H) and 7AI(D) tautomeric dimers to be 3.2 and 4.2 ns at 298 K and 13.3 and 16.9 ns at 77 K, respectively. Thus, the tautomeric dimer has a considerably longer lifetime than the tautomer in alcohols at room temperature. The dimers do show a lifetime ratio qI(D)/rfl(H) 1.3 that is very close to the one observed in alcohols. However, in the dimer case this isotopic ratio is independent of temperature. Another difference is that there is very little activation energy associated with the decay of tautomeric dimers in nonpolar solvents, unlike the case in alcohols. Based on these few comparisons it seems reasonable to conjecture that in 7AI/alcohol systems the nonradiative decay of the tautomer is related mainly to deactivation via N( I)-.H-OR interactions. This deactivation is similar to that operative in the 7AI dimer; however, it is considerably more effective in the alcohol case. Part of the difference could be the result of differences in the H-bonding properties of alcohols versus a partner tautomer. The relevance of such differences is indicated by the correlation
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(40) Waluk et aL2*measured these decay times over a wide temperature range and actually found that they were best fit to a function of the form In ( T ) = C + A exp(-E,/kT) having a nonzero value of C, rather than to a simple Anhenius form. Based on their Figure 3 it appears that the activation energy one would measure in the temperature region employed here would be smaller than the reported E, values and thus in closer agreement with the present 7AI results.
J. Phys. Chem. 1991, 95, 10369-10373 of Tfl with ET(30) (Figure 7b). The greater motional freedom of the pendant group in alcohols versus a dimer partner could also be the source of the difference. This latter explanation is favored by the fact that the lifetimes do correlate reasonably with viscosity (Figure 6b) and show an activated temperature dependence suggestive of the importance of large-amplitude motion in alcohols. IV. Conclusions
The results of the present study lead to a number of conclusions concerning the dynamics of excited-state proton transfer between 7AI and alcohol solvents. In addition to the proton-transfer reaction, time-dependent solvation of the initially excited normal species causes a time-dependent Stokes shift of the normal emission. The effect of this Stokes shift is apparent throughout most regions of the spectrum, complicating observation of the proton-transfer kinetics. The conflicting interpretations of the kinetics of 7AI/alcohol systems existing in the literature appear to have resulted from neglect of this previously unnoticed complication. Thus, after accounting for the improved signal-to-noise ratio afforded by the time-correlated single photon counting technique, the present results do not differ substantially from the primary data reported by others. It is merely the fact that different workers monitored emission at different wavelengths (often in regions where emission from the Stokes shifting normal band is significant) that has produced the varied interpretations reported for this reaction. The present results show that emission at wavelengths greater than 550 nm is essentially free from contamination by the normal emission and can be used to reliably measure the kinetics of the appearance of the tautomeric form. Results obtained in this way clearly show that the normal species is the direct kinetic precursor of the tautomer. In addition there may be some unresolvably fast tautomer formation, as has been previously suggested, but it accounts for a small percentage (
