J. Phys. Chem. 1985,89, 2355-2361
2355
Proton-Transfer Kinetics of 3-Hydroxyflavone: Solvent Eftects Andrew J. G. Strandjord and P. F. Barbara*+ Department of Chemistry, University of Minnesota, Minneapolis, Minnesota 55455 (Received: October 8, 1984)
Solvent effects on the excited-state intramolecular proton-transfer reaction of 3-hydroxyflavone have been investigated. In alcoholic solvents, the rate constant for the slow proton-transfer component kslowis observed to depend on the hydrogenbond-donating ability of the solvent. In strongly donating solvents, including 2,2,2-trifluoroethanol and 2-chloroethanol, k,,,, is significantly diminished as compared to the less donating alcohols, methanol and ethanol. The Arrhenius parameters for ksloware clearly dependent on the hydrogen-bond-donating ability of the solvent. The experiments also reveal that k,,,,, is not a strong function of solvent viscosity. A proton-transfer mechanism that can rationalize these observations is presented.
Introduction Laser studies on excited-state intramolecular proton transfer (ESIPT) are beginning to yield important new insight on the role that subtle hydrogen-bonding effects can play in small-barrier, ultrafast proton-transfer reactions in solution.l+ A particularly interesting and, to date, experimentally challenging example of ESIPT is found in the photophysics of 3-hydroxyflavone (3HF) and related molecules. Despite extensive research on the photophysics of 3HF,2-" many important aspects of the protontransfer mechanism remain obscure. The major obstacle has been the difficulty that has been encountered in attempting to characterize and understand the complex and subtle hydrogen-bonding effects that exist for this molecule in solvents such as alcohols and ethers. This aspect of the proton-transfer mechanism of 3 H F is related to many contemporary problems in the study of chemical reactions in interacting solvents.I2 The present understanding of the photophysics of 3 H F is due to the efforts of several research The actual protontransfer process was discovered by Sengupta and Kasha,2 who assigned the previously observed3 dual fluorescence bands of 3HF to a normal N * isomer (wavelength, maximum fluorescence = 408 nm) and tautomer T* isomer (wavelength, maximum fluorescence = 530 nm) which can be represented as The overall
-
Woolfe and Thi~tlethwaite~ and, independently, Itoh et aL5 showed that (i) the ESIPT process (N* T*) could be monitored by picosecond fluorescence spectroscopy and (ii) that the ESIPT dynamics are a strong function of solvent and temperature. Our group, Strandjord et a1.,6 demonstrated that the proton-transfer process has a fast and a slow kinetic component. We also showed that the fast and slow rate constants for proton transfer and kslow)could be measured independently by combining dynamic and static emission data. We suggested that the biexponential proton-transfer kinetics are due to two species (relaxed and unrelaxed forms of N*) that have different rate constants for proton transfer. In a separate study, Strandjdtd and Barbara,' reported a detailed study of hydrogen/deuterium kinetic isotope effects in the photophysics. We observed that the isotope effect on the slow kinetic component of the proton-transfer process was not a function of temperature, which is opposite to that usually observed for proton-transfer reactions (see Discussion). The final relevant references are due to McMorrow and Kasha* who recently showed that the proton-transfer kinetics of 3HF in alkane solvents are extremely rapid, even at cryogenic temperat u r e ~ . ~These ~ authors were able to show that the previous reports2ve of slow proton transfer in alkane solvents were in error due to trace amounts of hydrogen-bonding impurities. In contrast, the previous reports of slow proton transfer in hydrogen-bonding solvents, such as alcohols and ethers, are apparently c o r r e ~ t . ~ , ~ - ~ McMorrow and Kasha interpret the solvent trends by assuming that the proton-transfer barrier is zero (or at least very small) ~~~
Tautomer
Normal
excited-state proton-transfer process can be simply represented as hv
FT '
N-+N*+T* where N is the stable ground-state isomer and the superscript asterisk signifies that the denoted species occupies its lowest excited singlet state. Sengupta and Kasha have suggested that the T* species should be represented by the following canonical form:2
However, little information is available on the actual charge distribution of T*. 'Alfred P. Sloan Fellow and Presidential Young Investigator.
0022-3654/85/2089-2355$01.50/0
(1) (a) Huston, A.; Scott, G. S.; Gupta, A. J . Chem. Phys. 1982,76,4978. Bocian, D. F.; Huston, A. L.; Scott, G. W. J . Chem. Phys. 1983, 5802. (b) Nagaoka, S.;Hirota, N.; Sumitani, M.; Yoshihara, K. J . Am. Chem. SOC. 1983,105,4220. (c) Flom, S. R.; Barbara, P. F. Chem. Phys. Len. 1983,94, 488. (d) Ding, K.; Courtney, S. H.; Strandjord, A. J. G.; Flom, S. R.; Friedrich, D.; Barbara, P. F. J . Phys. Chem. 1983, 87, 1184. (e) Barbara, P. F.; Brus, L. E.; Rentzepis, P. M. J . Am. Chem. SOC.1980, 102, 5631. (f) Barbara, P. F.; Rentzepis, P. M. Brus, L. E. J . A m . Chem. SOC.1980, 102, 2786. (g) Goodman, J.; Brus, L. E. J. Am. Chem. SOC.1978, ZOO, 7472. (2) Sengupta, P. K.; Kasha, M. Chem. Phys. Lett. 1979, 68, 382. (3) Frolov, Y. L.; Sapozhnikov, Y. M.; Barev, S.S.; Pogodaeva, N. N.; Tyuhavkina, N. A. Izv. Akad. Navk SSSR, Ser. Khim 1974, 10, 2364. (4) Woolfe, G. J.; Thiitlethwaite, P. J. J. Am. Chem. Soc. 1981,103,6916. ( 5 ) Itoh, M.; Tokumura, K.; Tanimoto, Y.; Okada, Y.; Takeuchi, H.; Obi, K.: Tanaka. I. J . Am. Chem. SOC.1982. 104. 4146. '(6) Strandjord, A. J. G.; Courtney, S . H.;'Friedrich, D. M.; Barbara, P. F. J . Phys. Chem. 1983.87, 1125. (7) Strandjord, A. J. G.; Barbara, P. F. Chem. Phys. Lett. 1983, 98, 21. (8) McMorrow, D.; Kasha, M. J. Phys. Chem. 1984.88, 2235. McMorrow, D.; Kasha, M. J . Am. Chem. S0c.~1983,105, 5133. (9) Itoh, M.; Kurokawa, H. Chem. Phys. Lett. 1982, 91, 487. Itoh, M.; Tanimoto, Y.; Tokumura, K. J . Am. Chem. SOC.1983, 105, 3339. (10) Salman, 0. A.; Drickamer, H. G. J. Chem. Phys. 1981, 75, 572. Salman, 0. A.; Drickamer, H. G. J . Chem. Phys. 1982, 77, 3329. (1 1) Choi, K.;Boczer, B. P. In "Ultrafast Phenomena": Auston, D. H., Eisenthal, K. B., Eds.; Springer: New York, 1984, in press. (12) Coetzee, J. F.; Ritchie, C. D. "Solute-Solvent Interactions"; Marcel Dekker: New York, 1969. (13) It has been reported that 3HF exhibits a rapid but detectible proton-transfer rate at 77 Kin alkane solvents; see McMorrow, D.; Dzvgan, T.; Aartsma, T. J. Chem. Phys. Lett. 1984, 103, 492. The ESIPT rate was determined by monitoring the rise time of the T* fluorescence.
0 1985 American Chemical Society
2356
The Journal of Physical Chemistry, Vol. 89, No. 11, 1985
Strandjord and Barbara
TABLE I: Static and Dynamic Fluorescence and Absorption Parameters of 3HF in Various Solvents" p a x
solvent
cm-'
cm-'
fwhhNn, cm-'
methanol ethanol propanol butanol octanol 2-chloroethanol trifluoroethanol methylcyclohexane carbon tetrachloride
29 152 29 137 29 032 29 027 28 985 29015 29 154 29 498 29 188
24 630 24 727 24 69 1 24 752 24 691 24 175 24715
3 280 3 536 3 350 3 405 3 248 3 900 3 200
ah
3
YmN??,
v2;- p;, 4 490 4410 4 338 4 275 4 294 4 800 4 439
- urnax Til , cm-'
p a x abs
fwhhm, cm-'
U F ,
cm-
cm-' 18910
2001 1817 1960 1936 1967 2 I70 2 599 1540 1540
18 796 18 796
18 761 18 761
19047 20 040 19083 19047
10 229 10341 10232 10266 10224 9 953 9 114 10415 I O 141
"The error in these value is --f50 cm-'.
for unsolvated N* molecules-the only N* form that is possible in pure alkane solvents. This is similar to the behavior of many molecules that exhibit E S I P T . ' S ~In ~ contrast, the slow protontransfer kinetics in hydrogen-bonding solvents are ascribed to solute/solvent complexes which have a significant proton-transfer barrier. In spite of recent progress on the 3HF problem, a definitive quantitative description of the proton-transfer mechanism remains unavailable in the published literature. The problem lies in the lack of sufficient kinetic data, as a function of solvent and temperature, to evaluate the many mechanistic hypotheses2-11that have been put forth on the photophysics of 3HF. Perhaps the most important issue here is the physical basis of the slow protontransfer kinetic component. It has not been clearly established as to whether the slow component can be associated with activated barrier crossing in the sense of transition-state theory (tst), or whether the slow component is partially or wholly due to non-tst effects, such as (i) incomplete dielectric relaxation" of the solvent configuration, (ii) modulation of the isomerization rate due to coupling of the large amplitude motion of the solute to the configurational dynamics of the solvent,8 and/or (iii) quantum mechanical proton tunneling. The large amplitude motion effect is well-known in photophysics, especially in the isomerization dynamics of substituted ethylenes.16 It is often described as a solvent viscosity effect because when this mechanism is operating, the isomerization rate constant is observed to be a function of the solvent viscosity, q, which can be a measure of the solvent configurational dynarnics.l6 In an attempt to elucidate many of the remaining issues on the photophysics of 3 H F we undertook a detailed study of the photodynamics of this molecule. The present paper describes the effect of solvent variation on the photophysical dynamics. In the following paper in this journal, we present the results of a study of several structural derivatives of 3HF. The mechanistic implications of this work are discussed in depth with a particular emphasis on the physical basis of the proton-transfer kinetics. Experimental Section The picosecond and nanosecond fluorescence measurements described herein were made with an emission spectrometer that has been described in detail el~ewhere.~~'' The excitation wavelength for these studies is 355 nm. A Cary 17 spectrometer was employed in ultraviolet absorption measurements. The infrared data were recorded with a Beckman 4850 spectrometer. 3-Hydroxyflavone was purchased from Eastman Chemicals and purified by multiple recrystallizations from hexane. Sample so~~~~~~
(14) Rossetti, R.: Rayford, R.: Haddon, R. C.; Brus, L. E. J . Am. Chem. Soc. 1981,103,4303. Rossetti, R.; Brus, L. E. J . Chem. Phys. 1980, 73, 1546. Rossetti, R.; Haddon, R. C.; Brus, L. E. J . Am. Chem. SOC.1980, 102, 6913. Bondybey, V. E.; Haddon, R. C.; English, J. H. J . Chem. Phys. 1984, 80, 5432. (15) For a recent review of transition state theory, see: Truhlar, D. J.; Hase, W. L.; Hynes, J. T. J . Phys. Chem. 1983, 87, 2664. (16) Velsko, S. P.; Fleming, G. R. J . Phys. Chem. 1982, 76, 3553. Loufty, R. 0.;Arnold, A. J . Phys. Chem. 1982, 86, 4205. Evans, G. T.; Knauss, D. C. J . Chem. Phys. 1980, 72, 1504. (1 7) Barbara, P. F.; Strandjord, A. J. G. "Multichannel Image Detectors"; Talmi, Y . ,Ed.; American Chemical Society: Washington, DC, 1983; ACS Symp. Ser. No. 136, p 183.
Wavelength (nm)
Figure 1. Fluorescence spectra of 3-hydroxyflavone (5 X IO" M) at ambient temperature. The relative intensity of the different spectra have been adjusted arbitrarily. For these spectra, the excitation wavelength is at the absorption maximum; see Table I.
'/
b
300
'
"
"
325
\\\
TFE
'
'
f"
350
'~ 1 '
1
375
'
-
Wavelength (nm)
Figure 2. Absorption spectra of 3 H F at ambient temperature in methanol (MeOH), ethanol (EtOH), methylcyclohexone (MCH), 2-chloroethanol (2CE), and trifluoroethanol (TFE). The relative optical density of the different spectra has been adjusted arbitrarily.
lutions ( M) for the fluorescence measurements were made with solvents that were free of fluorescence impurities. The solvents, 2,2,2-trifluoroethanol, 2-chloroethanol, and ethanol, were freshly distilled. The quantity AvOH for the various alcohols was measured by the usual method.Is It was calculated from the formula A V O H = uHB - ufrec, where vfree is the frequency in cm-I of the O H stretching absorption of the alcohol under investigation (dissolved in CCl.,) and vHB is the frequency of the shifted OH stretch that appears when pyridine is added to the CCl,/ROH solution. Results Static Fluorescence and Absorption Experiments at Ambient Temperature. Dual emission bands are observed for 3 H F in a broad range of alcohol solvents as portrayed in Figure 1 and Table (18) Becker, E. D. Spectrochim. Acra 1961, 17, 436. Dureell, K. F.; Wilson, S. T. J . Mol. Spectrosc. 1967, 24, 468. Brandmuller, J.: Seevogel, K. Spectrochim. Acta 1964, 20, 453.
Kinetics of 3-Hydroxyflavone: Solvent Effects
The Journal of Physical Chemistry, Vol. 89, No. 11, 1985
2351
TABLE 11: Proton-Transfer Dynamics vs. Solvent Properties TB-l,
solvent methanol ethanol propanol butanol hexanol heptanol
octanol trifluoroethanol 2-chloroethanol
%
cp
0.59a 1.19 2.25 2.95 5.1 7.0 9.13 1.77 3 .OO
c
AvOH,cm-I
s-I x 10'0
32.4b 25.0 20.0 17.8 13.3 12.1 10.3 26.7 25.8
272c 257 257 257 256 258 260 418 325
2.86d 3.33 2.22 2.22 2.00 1.82 1.82 0.87 0.80
Ps 365d 400 550 475 655 880 975 500 825
FGIWBTG) 1 1.49 1.36 1.05 0.37 0.51
The viscosity values were found in the literature, as follows: Andrussov, L. "Zahlenwerte and Funktionen"; Landolt, H. H., Ed.; Springer-Verlag: West Berlin, 1950; I1 Band, S. Teil, p 1. bThe dielectric constant values were found in the literature as follows: (2CE), Sarojini, G. Ind. J . Pure Appl. Phys. 1972, 10, 526; (TFE), Schmid, R. J . Solution Chem. 1983, 12, 135; "Handbook of Chemistry and Physics"; Chemical Rubber Co.: Cleveland, OH, 1970. CSeeExperimental Section. dThe error of these values is &20%. They were determined by picosecond spectroscopy of the N* fluorescence.
I. In methylcyclohexane (MCH) and carbon tetrachloride (CCl,) (not shown in Figure l ) , however, only the T* fluorescence is observed. This indicates that the proton-transfer rate in M C H and CCl, is extremely rapid, which is apparently due to the fact that the proton-transfer rate for 3 H F in non-hydrogen bonding solvents ( M C H and CC14) is undiminished from its intrinsically rapid rate.8 Several spectroscopic parameters support the hypothesis that 3HF is hydrogen bonded to the solvent in alcoholic en~ironments.'~ The absorption spectrum of N is shifted toward lower energy as compared to 3 H F in MCH, as shown in Figure 2 and in Table is I, where the frequency of the first absorption maximum, ,$:v listed for various solvents. Further evidence is found in the emission spectra, for which the most strongly hydrogen-bonding solvents, 2-chloroethanol (2CE) and 2,2,2-trifluoroethanol (TFE), exhibit some significant differences from the unsubstituted linear alcohols (see Figure 2 and Table I). In particular, 2CE has the lowest energy N* emission maximum (vAaflX), the largest N* fluorescence Stokes shift (vzt- vItfanx), and the largest full width half-height of the N* emission (fwhhNfl). An even larger effect is observed for the T* emission of 3 H F in TFE. Compared to the other alcohols and MCH, the emission maximum of T* (vTmflax) is shifted to greater energies by 1000 cm-l or more in TFE. In summary, the static spectroscopic parameters of 3 H F in a broad range of solvents support the hypotheses that the N and T forms of 3 H F are hydrogen bonded in alcoholic solvents and that the energy of N, N*, T, and T* is modulated by hydrogen-bonding interactions. Definition of the Photodynamical Observables. A thorough investigation of the excited-state proton-transfer dynamics of 3HF requires both static (Figure 1) and dynamic (Figure 3) fluorescence data.7-20 The single necessary observable from the static data is the ratio of intensities of the green (T*)to blue (N*) emission bands, which is denoted here by FGIFB.Several parameters, however, are extracted by a fitting procedure from the fluorescence dynamics6 The N* fluorescence decays with two time constants, which are associated with the slow and fast components of the proton-transfer mechanism. Accordingly, the T* emission dynamics exhibits two formation time constants. In many environments the shorter of the two time constants is much shorter than the resolution of our apparatus (-30 ps). Under these circumstances, the N* emission appears to exhibit only the longer time constant, which we denote in this paper as 7B, the lifetime of the blue (N*) emission. In turn, the T* fluorescence formation kinetics appears to exhibit an "immediate" and "delayed" component as shown in Figure 3b. The T* kinetic traces can be analyzed to determine the ratio of fast (immediate) to slow (delayed) quantum yields of the T* (green) emission. This ratio is denoted by &jfast/&dow. Another parameter that is necessary for determining the proton-transfer (19) Matoga, N.; Kubota, T. "Hydrogen Bonding Complexes"; Marcel1 Dekker: New York, 1970; Chapter 7, p 293. (20) In ref 6 and 7 Fo/(FeTG) is denoted by F ~ / ( ~ Tand c )q5gf*at/&ci'oW is represented by ~ C f a s t / ~ C a ' " w .
I
I
I
0
200
400
I
Time (psec) Figure 3. Time-resolved emission traces at 405 nm (a) and 530 nm (b) for 3HF in methanol (0)and 3HF in methanol-dl ( 0 )induced by 355nm 30-ps laser pulses at 240 K. The points in the figure (0 or 0) represent experimental data. The solid lines are computer fits of the data that take into account the time response of our apparatus; see Strandjord and Barbara.'
dynamics is T G , which is the lifetime of the T* (green) fluorescenceS6 Proton- Transfer Kinetics at Ambient Temperature. The ex* an especially useful parameter in the perimental quantity T ~ - is investigation of the proton-transfer dynamics at ambient temperature. In general, 7B-l is the sum of all radiative and nonradiative rate constants responsible for the disappearance of N* (see Appendix). In alcoholic solvents near ambient temperature, however, the proton transfer is more rapid than the other processes, and, therefore, T ~ is- approximately ~ equal to the slow protontransfer rate constant, kslow.'l Table I1 lists T ~ - ' for solutions of 3HF in various straight-chain alcohols. The proton-transfer dynamics are significantly slower in the more strongly hydrogen-bond-donating alcohols T F E and 2CE. The quantity AuoH is a known measure of the hydrogenbond-donating ability of alcohols.'8
2358
Strandjord and Barbara
The Journal of Physical Chemistry, Vol. 89, No. 11, 1985
CI
V
-
> -
2.0
1.0
t I-'
i
-
2CE
2
{ , ,
I
4
6
I
, ,
TABLE 111: Proton-Transfer Activation Parameters
fractiona Emb slow (fs) f . F c / ( F n ~ c ) kcal/mol 0.58 0.59 0.61 0.53 0.74
1 .oo 1.52 0.60 0.34 0.54
1.4 1.7 5.9 4.5 1.4
A: s-' 1.77 X 10l2 2.75 X 10l2 1.6 X 10l5 1.0 x 1014 6.16 X 10"
"The error of these values is -*0.05. bCalculated from the data shown in Figure 5, error -*0.1 kcal/mol. 'See text for a description of this quantity
In addition to the hydrogen-bonding solvent effect on ~ g - l a, mild viscosity 7 effect is also apparent. The viscosity effect is illustrated in Figure 4 by a plot of k,,,, (re-') vs. 7 for the unsubstituted linear alcohols. Interestingly, ksl, for the more strongly interacting solvents T F E and 2CE does not fall on the same line as the unsubstituted alcohols. This strongly suggests that the hydrogen-bonding effect dominates the viscosity effect for the solvents like TFE and 2CE in the viscosity range of Figure 4. Compared to many examples of viscosity effects on isomerization rates of electronically excited moleculesx6the 3 H F trend is quite mild. For example, consider the torsional motion of tetraphenylethylene2' for which ktorSIOn a 7-l. In contrast ( rB-') has a finite zero viscosity intercept k(q = 0) N 3 X loxos-l. This implies that near ambient temperature kslowis not due entirely, or even predominately, to viscosity effects on the rate of solvent configurational motion. Instead, energetic factors involving the 3HF/solvent hydrogen bond seem to be the dominating parameters controlling kslow. This interesting observation will be discussed in detail below (see Discussion). An alternative measure of kslowcan be found in eq 1 (see k i O w=
( k r ~ / k r ~ ) f , F ~ / ( F ~ 7 ~ ) (1)
Appendix). Here krNand krTare the radiative rate constants of the N* and T* forms, respectively, and f s is the fraction of the slow yield of proton transfer as defined by eq 2. Table I11 lists fs = 4Gfast/(4Gfast
+ dGslow 1
4
5
1/Tx 10'
Viscosity Figure 4. Plot of proton-transfer rate constant (re-') for 3HF vs. solvent viscosity of various linear alcohols at room temperature; see Table 11.
solvent methanol ethanol 2-chloroethanol trifluoroethanol methanol-d,
3
8
(2)
relative experimental values fdr the right side of the eq 12 as a function of solvent. These values can be used as a relative measure of k,,,, if it is assumed that the radiative rate constants are not significantly solvent dependent. The solvent dependence of this measure of kslow is similar to that previously mentioned for rB-l in Table 11. The measurements in Table 111, however, are less reliable due to the large experimental uncertainties in the f,,,, values. (21) Vinogradov, S.N.; Linnel, R. H. "Hydrogen Bonding"; Van Nostrand Reinhold: New York, 1971.
Figure 5. Arrhenius plots of log FG/(FBTG) for 3HF in several alcoholic
solvents. The limited data in Table I11 seem to suggest thatf, is similar for alcohols of very different hydrogen-bonding ability. This allows one to use the experimental quantity FG/(FBTG) as a relative measure of k,,,. As expected, a similar solvent trend is observed for this quantity to that found for rB1I. Temperature Dependence of the Proton- Transfer Kinetics. We are particularly interested here in determining the temperature dependence of kslowwhich, in principle, can be extracted from Arrhenius plots of the right side of eq 1. However, the experimental determination of the quantity f s is especially difficult and time consuming. We have limited our observations, therefore, to the more easily determined complex observable FG/(FBTG). Previous work from our laboratory suggests thatfs is only weakly temperature dependent in alcoholic solvents.' Consequently, the data in Figure 5 should offer a good measure of the Arrhenius energy E, of kslow(Table 111). It is particularly interesting that E , is significantly greater for the more strongly hydrogen-bonding solvents, TFE and 2CE. The Arrhenius A factor is also clearly dependent on solvent (Table 111). The A values were determined by combining the data in Figure 3 with the ambient temperature values for 7B-I which, as stated previously, are approximately equal to kslow.
Discussion The major goal of this work is to gain new insight into the "mechanism" of FSIPT of 3HF in alcohols. This is made difficult by the inherent complexity of ultrafast proton-transfer processes in hydrogen-bonding environments. One anticipates many possible modes of dynamic and static coupling of the solvent and reacting solute in reactions of this type. It may be that a truly detailed understanding of 3HF solute/solvent dynamics will not be attained until molecular dynamic simulations of this solute/solvent system are accomplished. Nevertheless, the large amount of kinetic data now available on the photodynamics of 3HF encourages us to formulate and evaluate several hypothetical schemes for the FSIPT mechanism. In constructing these models, we have liberally incorporated the mechanistic ideas of many of the workers in this field, especially Choi et a].," McMorrow and Kasha,* and Thistlethwaite and W001fe.~ Physical Basis of kslow. The results presented in this paper strongly suggest that the slow proton-transfer mechanism should be viewed as a thermally activated barrier crossing, or at least a mechanism closely related to this type of a process, such as a thermally activated quantum mechanical proton tunneling process. This conclusion can be drawn from a variety of evidence. First of all, kslowexhibits an Arrhenius dependence that is consistent with a thermally activated atom-transfer reaction in solution; Le., the plots are linear and have an A factor in the range 1010-1014 SKI.Second, kslow seems to depend on energetic factors, rather than solvent relaxation effects, as evidenced by the absence of a strong k,,,, vs. solvent viscosity dependence. Finally, experiments on the photodynamics of several derivations of 3HF, which are described in the following article in this journal, are highly sup-
The Journal of Physical Chemistry, Vol. 89, No. 11, 1985 2359
Kinetics of 3-Hydroxyflavone: Solvent Effects
o/
H
I
H\
R
,0-H R
0-R
H\
R-0"
0-R
I R
/
0
I
R
H... \
\
0
0
R
R
I
I
T*
NY*
portive of the thermal activation hypothesis. Another conclusion that can be drawn from the results of this paper and previous work is that the reaction coordinate that is involved in the thermal activation of kslowmust at least partly be due to motion of atoms associated with an intermolecular hydrogen bond between N and the solvent. The major evidence for this hypothesis is the apparent positive correlation we observe between E, and the strength of the intermolecular hydrogen bond which can be inferred qualitatively from the hydrogen-bond-donating ability of the solvent. The lack of a significant observed correlation between viscosity and k,, tend to discount hypothetical proton-transfer mechanisms that are based upon large amplitude motion of the solvent. It is unlikely, for example, that the correct mechanism for ESIPT resembles a sequence as shown in eq 3. The desolvation step in eq 3 would be expected to be sensitive to the configurational dynamics of the solvent which in turn should be manifested by the solvent viscosity, 7. For a hypothetical scheme like that shown in eq 3, we would expect that kslowwould be experimentally correlated with the solvent viscosity. Dynamical Mechanism Based on a SolventlSolute Complex. W e now discuss a mechanistic picture that can rationalize the broad range of experimental data on the photodynamics of 3 H F in alcohols. The mechanism is derived from the simple premise that the proton transfer of 3 H F in alcohols should be viewed as an intramolecular proton transfer of a hydrogen-bonded solute/solvent 3HF- - -HO-R complex. The model is best discussed by considering some hypothetical structures for the solvated forms of N* and T*. Our primary goal here, however, is to demonstrate the salient features of dynamical aspects of the mechanistic picture rather than to establish the actual structure of the N* and T* 3HF- -ROH complexes. A particularly attractive scheme for the proton transfer can be represented as in eq 4. In this equation N, represents the equilibrium (reactant) geometry of N*, while N,* is a thermally excited form of N,* that occurs along the reaction coordinate for the N* T* proton transfer. N,* cannot directly proton transfer because its intramolecular hydrogen bond is weakened by the presence of a strong intermolecular hydrogen bond. Ny*,on the other hand, undergoes rapid proton transfer because its intermolecular hydrogen bond is almost broken. In other words, Ny* resembles an unsolvated N* molecule. As such, Ny*is capable of very rapid proton transfer (see Introduction). The underlying theme in eq 4 is that the ESIPT process in alcohols occurs by thermal activation in the transition-state-theory sense, perhaps with a significant quantum-mechanical-tunneling component. According to this picture, the alcohol kinetic effect
-
Model E I S o l v a t i o n II-
P
r
I
Proton Transfer
u/ T'
Figure 6. Schematic representation of hypothetical potential energy curves for the lowest excited singlet state for Model A and Model B; see text for further details.
should be classified as a "static" solvent effect,15since the kinetics of the actual proton-transfer process is not a function of relaxation times associated with the solvent configuration. Rather, the proton-transfer rate constant depends on the energetics of hydrogen bonding. Structure of the Direct Precursor to ESIPT. The physical significance of the model reflected by eq 4 can be further developed by making a detailed kinetic analysis of the photodynamic implications of this mechanism. For the sake of discussion, we consider two limits to the model implied by eq 4. The limits are outlined in Figure 6. In model A the transition state is assumed to occur before the actual proton-transfer process. The participation of motion along coordinates involving the intermolecular hydrogen bond is portrayed in Figure 6 as "solvation". Model
2360 The Journal of Physical Chemistry, Vol. 89, No. 11, 1985 A is a two-step mechanism. The first step is thermally activated desolvation, followed by rapid proton transfer. The position of the transition state in model B is further along the reaction coordinate. In model B the rate-limiting step to overall proton transfer involves the simple proton-transfer process. As a result, quantum mechanical proton tunneling could play a key kinetic role in model B, but would be less important in model A, for which the transition state occurs along coordinates involving solvation. Model B, incidently, resembles the underlying energetics of the Marcus theory of proton transfer.22 Before we can evaluate whether model A or model B more accurately describes the actual proton-transfer mechanism, it is necessary to develop reasonable kinetic schemes for photodynamical manifestations of these models. In constructing the kinetic schemes we have attempted to realistically model the protontransfer mechanism implied by Figure 6. We have also included the important features specifically associated with the photodynamics, such as the radiative and nonradiative behavior of N* and T*. In model A the rate-limiting step to the observed proton transfer is motion along coordinates involving solvation (Figure 6). The fast component is due to direct production of Ny* (by optical excitation of N,) followed by unreasolvably rapid proton transfer (N,* Ty*). The slow component in this model is due to the sequence N,* Ny* T*
-
- -
The parameters of model A can be related to the experimental observables by eq 5-8, the derivation of which is given in the kslow = k2 7g-l
+ knrN+ k2
(6)
= Kgsey/(ex7Bk2)
(7)
= krN
d'fast/d'slow
(5)
Kgs = kl/k-l
(8)
Appendix. Here ex and e, are the molar extinction coefficients of N, and N,, krNand krTare the radiative rate constants for N, and N,, and knrNand knrTare the sum of all nonradiative rate constants (except k,) for N,* and T* (see Figure 7 ) . In Model B the proton-transfer dynamics (eq 9) depends on N, e N, T* (9)
-
both the solvation coordinates and the microscopic proton-transfer process. An important feature of model B is that it includes excited-state relaxation of the N forms: Ny*
k-2
N,*
This relaxation is the essential element of a model we had previously proposed. However, model B also includes the ground-state equilibrium (N, e N,), which was missing from our previous model, but now seems to be justified based on the arguments of McMorrow and Kashas and Choi et a1.li A kinetic scheme for model B is outlined in Figure 7 . The fast proton transfer is associated in model B with direct excitation of N,, in an analogous fashion to model A. The physical significance of model B is that N,* competes is assumes that the rate of the relaxation N,* with the microscopic proton-transfer rate. The kinetic expressions for model B are shown in eq 10-13 in -+
ksiow = k2(1 - Y) TB-'
= krN + knrN + k2(l - Y) Y = k-2/(k-2
d'fast/d'slow
+ k,)
= KgseyrB-'/ [ M e , + Y&sey)I
(10) (11) (12) (13)
which the quantities krN,krT,e,, e,, and KBcare defined in the same way as they were for eq 5-8. (22) Marcus, R. A. J . Phys. Chem. 1968, 72,891. Cohen, A. 0.;Marcus, R . A. J . Phys. Chem. 1968, 72, 4249.
Strandjord and Barbara
I
'-
,xNy1
Nx
N"
T
Figure 7. Photodynamic scheme for model A and model B. It is assumed that k3 >> k2, kT2for model A.
Model A us. Model 8. The essential difference between the two models is that for model A motion along solvation coordinates is rate-limiting (eq 5 ) , while for model B both solvation coordinates and proton-transfer coordinates participate in the slow component kinetics (eq 10 and 12). The question as to which point of view is more valid, Le., that implied by model A or that implied by model B, has been an underlying theme in the photophysics of 3HF. Fortunately, it now appears that there is sufficient data to answer this question-at least for alcoholic environments. The revealing parameter here is the hydrogendeuterium isotope effect for kslow,which can be measured by the quantity kslow= f,FG/(FBrG). We previously observed' that this measure of kslOw shows an isotope effect when compared the solvents MeOH and MeOD, Le., ksl,(MeOH)/k~low(MeOD) = 2.2 f 0.03. In MeOD the alcoholic proton of 3HF is, of course, exchanged by a deuteron. If this isotope effect is a manifestation of the microscopic pro! T*), then model A is cast in doubt, ton-transfer proces (Ny*& since kslowis not a function of k3 in model A. Unfortunately, there is some ambiguity as to the identity of the process that is responsible for isotope effect. The problem is that N,* has two labile protons, i.e., the one that is transferring and the alcoholic proton in ROH. Presumably, both sites are exchanged with a deuteron in the solvent MeOD. The observed isotope effect may, therefore, be due to either k,, k3, or both. In our estimation, the most likely candidate, however, is k,, the microscopic proton-transfer rate constant. We base this conjecture on the close similarity of kH/kDin methanol (2.2 f 0.03) and kH/kDin ethyl ether (2.35 f O.2).' In ethyl ether the solvent is not deuterated and only the transferring proton is exchanged in the isotope studies. Although the proton-transfer mechanism in ethyl ether (a hydrogen-bond-accepting solvent) may be quite different than model A and model B, the similar isotope effects suggest that the same microscopic proton-transfer process may be operating in both types of solvents. Additional evidence that disputes model A is found by examining the isotope effects on the quantity ~ f a s t / ~ s l o wwhich can be calculated from thef, values in Table 111. The result is (dfast/ d'slow)McOH!(d'fast/~slow)McOD = 2.18 at ambient temperature. To interpret this result, we should consider the right side of eq 7. Since iB-'= k2 at ambient temperature (in model A), eq 7 can be approximated by eq 14. Equation 14 shows that, if model A is d'faat/4Jalow = Kgaey/ex (14) valid, the isotope effect we find for I#Jfast/I#Jslow must be due to the equilibrium constant of ground-state normal forms ( K , = k,/k-,) and/or the molar extinction coefficient of N, and N,. It is unlikely,
The Journal of Physical Chemistry, Vol. 89, No. 11. 1985 2361
Kinetics of 3-Hydroxyflavone: Solvent Effects however, that ex, cy, or Kgs should be significantly sensitive to isotopic substitution. The absorption spectrum of 3HF in MeOH and the spectrum in MeOD are identical within experimental error, which casts doubt on the possibility of isotope effects on cx and cy. The expected lack of a significant isotope effect on Kkris based on analogy to the usual behavior of hydrogen-bonding equilibrium constants in alcoholic solvents.21 In conclusion, an interpretation of two types of isotope effects casts doubt on the validity of model A. On the other hand, the isotope effects are highly consistent with model B, for which an assumed isotope effect on the microscopic rate constant k3 would explain the experimentally observed effect. Some insight into the proton-transfer process in model B can be gained from a consideration of the hydrogen/deuterium isotope data on kslowand the experimental variable, ~fast/&,w. The lack of a significant temperature dependence of the isotope effectsI2 may be an indication that quantum mechanical tunneling is the dominant mechanism for proton transfer. Clear evidence for tunneling has, in fact, been observed for a variety of excited-state intramolecular proton-transfer molecules, including the widely studied methyl salicylate'g and other compound^.'^ Conclusions and Summary Solvent effects on the proton-transfer kinetics of 3-hydroxyflavone have been studied. In alcoholic solvents it is observed that the thermal activation parameters (Ea and A) of the slow proton-transfer rate constant (kslow)depend on the hydrogen-bonddonating ability (AvoH) of the solvent. The experiments also reveal that k,,,, is not a strong function of the solvent viscosity. The observed behavior can be rationalized by a phenomenologically oriented mechanism (model B) that assumes that 3 H F is complexed with the solvent. The temperature dependence of ksIwin this model is due to thermal excitation of coordinates that are associated with the intermolecular hydrogen bond of a solute/solvent complex. It is assumed that the thermally excited solute/solvent complex is the direct precursor to proton transfer. In the following companion article, we examine the effect of structural variation on the proton-transfer kinetics of 3hydroxyflavonols. The data are employed to examine the validity of the mechanistic hypotheses that are presented herein. The companion article also includes a discussion of the physical basis of the microscopic proton-transfer process of 3HF.
Nx*(t = 0) = Nxc,C
(A4)
Ny*(t = 0) = 0
('45)
P(t = 0) = NycyC
('46)
equilibrium ground-state concentrations, and C is a proportionality constant which takes into account the experimental conditions, e.g., excitation light intensity. It is important to note that Ny*(t = 0), at steady-state, is assumed to be equal to zero. Correspondingly, P(t = 0) during the steady state is equated with NyeyC,the actual initial concentration of Ny*. In other words, the initial concentration of Ny* is rapidly dissipated by the process Ny* T*, as the steady-state regime is established. Solutions to eq A4-A6 are given by the following equations.
-
Tslaw*(t)
Nx*(t) = Nx*(0) exp(-t/TB)
('47)
Ny*(t) = Ny*(O)(k,/k3)NXcxCexp(-t/TB-')
(AS)
T*(t) = Tfast* + Tslow*
(A9)
Tfast*(t) = CNyey exp(-t/TB)
('410)
= CNxcxk2TBTG/(TG - 7B)[exP(-t/TG) - exp(-t/TB)l (A1 1)
Equations 1 and 4-7 can be derived for eq A7-AlO by simple manipulations employing the definitions: FG
OC
FN
krTJmr*(t) dt
('412)
krNJmN*(t) dt
('413)
The derivation of eq 8-1 1, which are associated with model B, can be accomplished in a similar fashion to that just outlined. The appropriate differential equations are given by dN,*/dt = k-2Ny* - k2NX*
+ k3)Ny* = k3Ny* - (krT + k,,T)T*
dNy*/dt = k2NX*- (k-2 dP/dt
The initial conditions for photoexcitation in model B are given by eq A14-Al6. Here Y is given by eq 10.
Acknowledgment. Acknowledgment is made for partial support of this research to the donors of the Petroleum Research Fund, administered by the American Chemical Society, and the National Science Foundation (CHE-8351158). P.F.B. thanks Professor Maurice M. Kreevoy and Dr. K. Choi for helpful discussions.
Nx*(0) = CNxc, + CYNycy
(A14)
N,*(O) = 0
(A15)
P ( 0 ) = C(l - Y)Nycy
('416)
Appendix The derivation of eq 1 and 5-13 are outlined in this section. The differential equations corresponding to model A are given by eq Al-A3. The symbols are defined in Figure 7 and the dNx*/dt = -(krN + knrN k2)Nx* ('41)
The solution for the time-dependent concentration in model B are shown in eq A17-A21. Here TB = krN+ knrN+ k2 (1 - r). Ny*(t) = [k2/(k-2 + k3)lNx*(t)
(A181
dNy*/dt = k2Nx* - k3Ny*
('42)
T*(t) = Tfast*(t) + Ts~ow*(t)
(A19)
dr*/dt = k3Ny* - (krT + knrT)T*
('43)
+
Discusion section. Here it has been assumed that k3 > krN, knrN. The solution to eq Al-A3 is simplified if we make the steady-state assumption for Ny*. This is appropriate for model A, because the rate-limiting step to proton transfer is N,* -* Ny*. The initial conditions that apply for the case of pulsed photoexcitation at t = 0 are given by eq A4-A6. N, and Ny are the
Nx*(t) = [cNx(o)(cx + yKgscY)l exp(-t/7B)
Tfast*(t) = cyKg~Nx(0)~Y exp(-t/TB) Tslow*(t) = CYk2Nx(0)~B-1(ex + Yk,,cy) [exp(-t/rG)
(A17)
(A20)
- exp(-t/rB)] (A21)
Equations 1 and 8-1 1 follow in a straightforward fashion from eq A l l , A12, and Al6-A20. Registry No. 3HF, 577-85-5.
