J. Phys. Chem. 1996, 100, 16053-16057
16053
Infrared Spectroscopy of Hydrogen-Bonded Phenol-Amine Clusters in Supersonic Jets Atsushi Iwasaki, Asuka Fujii, Takeshi Watanabe, Takayuki Ebata, and Naohiko Mikami* Department of Chemistry, Graduate School of Science, Tohoku UniVersity, Sendai 980-77, Japan ReceiVed: March 7, 1996; In Final Form: June 25, 1996X
We report infrared spectra of hydrogen-bonded phenol-amine clusters, phenol-NH3, -N(CH3)3, -NH(C2H5)2, and -N(C2H5)3, prepared in jet expansions. The OH, NH, and CH stretching fundamentals were studied. Infrared-ultraviolet double-resonance techniques were utilized for vibrational spectroscopy of sizeselected clusters. The OH stretch frequencies of the phenol moieties showed extremely large red-shifts from that of bare phenol, reflecting the strong proton affinities of the amines. Moreover, non-proton-transferred structures of the clusters were confirmed. The detailed structure of phenol-NH3 was examined by ab initio calculations, which reproduced the observed infrared spectrum.
Introduction Recently, vibrational spectroscopic studies of low density molecular clusters produced in supersonic jets have become popular for their cluster structure analyses. In these studies, double-resonance techniques for highly sensitive ionization or fluorescence detection of infrared1-6 (or stimulated Raman2,7) transitions have been utilized. These techniques have also allowed us to select the cluster size of interest. OH stretching vibrations of hydrogen-bonded clusters, such as phenol(H2O)n,1,3,4 phenol-dimer,2 -trimer,2 and benzene-(H2O)n5,6 have been reported so far, and their cluster structures have been analyzed based on their vibrational spectra. Changes in the OH vibrational frequency as well as the intensity of the OH fundamental upon cluster formation have been found to be quite sensitive to cluster structure. In particular, characteristic frequency shifts have provided us with direct information as to intermolecular hydrogen bonds with various solvent molecules. The analyses have often been performed with the aid of ab initio calculations of cluster structures and of spectral simulations based on the energy-optimized structures.4,8-11 Comparison between the observed and calculated spectra of OH vibrations is one of the most reliable tests for cluster structures. In this paper, we report infrared (IR) spectra of hydrogenbonded clusters of phenol with various amines, such as NH3, N(CH3)3, NH(C2H5)2, and N(C2H5)3. The OH, NH, and CH stretching vibrational region is observed. Amines are much stronger bases than water, and hydrogen bonds between amines and phenol can be regarded as prototypes of strong hydrogen bonds. Though NH3 is the weakest base among the amines studied in this work, we find that the OH stretch vibrational frequency of phenol-NH3 shows a significant frequency reduction. The structure of phenol-NH3 is examined by ab initio calculations, which reproduces the observed infrared spectrum very well. For a hydrogen-bonded cluster involving a proton acceptor with large proton affinity (PA), intracluster proton transfer may be expected,8-15 as has been found in the S1 state of phenol(NH3)n with n > 4, for example.12,14,17,18 Since the alkylsubstitution of ammonia results in a large increase of PA, it is interesting to investigate whether intracluster proton transfer does occur even in a ground state of the hydrogen-bonded cluster with amine. The investigation can be performed by observation of the OH stretching vibrations of the phenol moiety. The X
Abstract published in AdVance ACS Abstracts, September 1, 1996.
S0022-3654(96)00711-3 CCC: $12.00
results in the present work show that intracluster proton transfer is not induced even in the cluster with the amine of the strongest basicity, though the OH frequency is reduced by 600 cm-1 from that of bare phenol. The frequency reduction is discussed with respect to the strength of hydrogen bonds. Experimental Section Hydrogen-bonded clusters were produced by supersonic jet expansions of a gaseous mixture of phenol, amine, and He into a vacuum chamber. The stagnation pressure of the gaseous mixture was 3 atm, where vapor pressure of amine was maintained by controlling the temperature of a compartment of amine. A typical background pressure of the vacuum chamber was 1 × 10-5 Torr. Details of IR spectroscopy using double-resonance techniques were previously described elsewhere.1 Only a brief description is given in the following. Clusters in jet expansions were resonantly ionized by a pulsed ultraviolet (UV) light beam whose wavelength was fixed at the 0-0 band of a particular cluster. The produced ions were mass-selected by a time-offlight mass spectrometer and were detected by a channel electron multiplier (Murata). A pulsed IR light beam was introduced prior to the UV beam, and its wavelength was scanned. When the IR wavelength was resonant with a vibrational transition of the cluster, the reduction of the ground state population was induced and was detected as a decrease of the ion current intensity. Hereafter, we call infrared spectra observed by this method ionization detected infrared (IDIR) spectra. The tunable IR beam was generated by difference frequency generation with a LiNbO3 crystal. A second harmonic of a Nd: YAG laser (Quantel YG581-10) and an output of a dye laser (Quantel TDL50, DCM dye) pumped by the YAG laser were mixed in the crystal. Typical power of the IR beam was 0.2 mJ, and its spectral resolution was about 1 cm-1. The beam was focused into the vacuum chamber by a CaF2 lens of f ) 250 mm. The UV laser beam was a second harmonic of an output of a dye laser (Lambda Physik FL3002, C540A dye) pumped by a XeCl excimer laser (Lambda Physik, LPX100). The UV laser power was typically 10 µJ, and its resolution was about 0.3 cm-1. The UV beam was focused by a f ) 500 mm lens and was counterpropagated with the IR beam. The delay time between the IR and UV pulses was set to be 50 ns with a digital delay generator (BNC), the former being prior to the latter. © 1996 American Chemical Society
16054 J. Phys. Chem., Vol. 100, No. 40, 1996
Figure 1. Ionization detected infrared (IDIR) spectra of (a) bare phenol, (b) phenol-H2O, and (c) phenol-NH3.
Iwasaki et al.
Figure 2. (a) IDIR spectrum of phenol-NH3, of which the ordinate was converted to absorbance. (b) Calculated infrared spectrum for the energy-optimized structure of phenol-NH3. The histogram shows the relative infrared intensity.
Results and Discussion Figure 1 shows the ionization detected infrared (IDIR) spectra of (a) bare phenol, (b) phenol-H2O, and (c) phenol-NH3. The spectra a and b are the same as those reported in a previous paper.1 Since the S1-S0 electronic transitions of these clusters have been well studied,12,13,20-22 the IDIR spectrum was obtained by tuning the UV laser wavelength at the 0-0 band of each cluster. 1. Phenol-NH3. The spectrum of phenol-NH3 is shown in Figure 1c, in which an intense band and a weak band are seen at 3294 and 3333 cm-1, respectively. Two weak bands also appear at 3058 and 3088 cm-1, where CH stretching vibrations are expected. Since OH stretching vibrations are known to be generally intense in IR spectra, the intense band at 3294 cm-1 can readily be assigned to the OH oscillator of phenol. On the other hand, the weak band at 3333 cm-1 can be assigned to an NH stretching vibration, which appears generally to be weak in IR spectra. These assignments are readily confirmed by the simulated spectrum by ab initio calculations of phenol-NH3, as described below. The IR spectrum obtained in this work partly corresponds to the Raman spectrum reported by Hartland et al.,7 who observed phenolNH3, as well as various other clusters by means of ionizationloss stimulated Raman (ILSR) spectroscopy. In the ILSR spectrum, only the NH band (3334 cm-1) was observed, but the OH band was not reported. The OH stretching frequency exhibits a large frequency redshift of 363 cm-1 from that of bare phenol (Figure 1a). The red-shift is extremely large, being compared to a red-shift of 133 cm-1 observed in phenol-H2O (Figure 1b). The NH band, on the other hand, shows a small red-shift of only 4 cm-1 from ν1 (3337 cm-1)24 of bare NH3. Another NH band of phenolNH3 corresponding to ν3 (3444 cm-1)24 of bare NH3, which is also IR active, was not found in the present spectrum. The large red-shift of the OH band and the small shift of the NH band represent that the phenol site is acting as a proton donor and NH3 as an acceptor and that the force field of the OH bond is reduced by the hydrogen bond formation, but that those of the NH bonds are almost unchanged. To analyze the cluster form on the basis of the IR spectrum, ab initio HF/SCF/MO calculations of an energy-optimized cluster structure and simulation of IR transitions were carried out by using the Gaussian 92 program.24 In Figure 2b, the calculated band positions and transition intensities obtained at the HF/6-31G level are illustrated as histograms. The observed spectrum of phenol-NH3 is reproduced in Figure 2a with an expanded energy scale for the abscissa and with absorbance for the ordinate, so that the intensities of the observed bands can be compared with the calculated. The details of the conversion method of the ion dip intensity to the absorbance
Figure 3. Side and top views of the energy-optimized structure of phenol-NH3 (6-31G basis) and definitions of the structural parameters that were listed in Table 2.
were given elsewhere.2 The simulated spectrum well reproduces the observed one, and the assignments of the bands described above were confirmed. The calculation presents an energyoptimized cluster structure and its harmonic frequencies; the side and top views of the energy-optimized structure of phenolNH3 are shown in Figure 3. Assignments of the observed IR bands are given in Table 1, being performed by a comparison between the scaled frequencies and the observed. The scaling factor for the calculated frequencies was 0.8997. This value was obtained from a linear least-squares fitting of the calculated and the observed frequencies of the CH, OH, and NH bands. Some vibrational frequencies observed in dispersed fluorescence and ILSR spectroscopic studies7,11,22 are also listed in Table 1, as well as those from the present IDIR study. The calculated frequencies of the intermolecular vibrations and other low-frequency vibrations of the phenyl ring also agree well with the observed, though the scaling factor was determined by fitting for the high-frequency vibrations. The consistency with the observed vibrational frequencies indicates that the calculated structure shown in Figure 3 is highly reliable. Table 2 shows the structural parameters of the energy-optimized structure of phenol-NH3. Those of bare phenol and phenolH2O calculated at the same level are also listed in the table. The OH bond distance (r2) of the phenol moiety in phenolNH3 is much longer than that of bare phenol. On the other hand, that distance in phenol-H2O is definitely smaller than in phenol-NH3. This is consistent with the hydrogen bond in phenol-NH3 being much stronger than that in phenol-H2O, reflecting the larger proton affinity of NH3 (204 kcal/mol) than H2O (116.5 kcal/mol).25 Recently, Schiefke et al.11 reported the dispersed fluorescence, two-color resonance multiphoton ionization, and hole-burning spectra of phenol-NH3. They also carried out ab initio calculations providing an energy-optimized structure, of which geometrical parameters are essentially reproduced in the present
IR Spectroscopy of Hydrogen-Bonded Phenol-Amine Clusters TABLE 1: Assignments of Calculated Harmonic Frequencies and IR Intensities of Phenol-NH3 Compared with Observed Frequencies (Calculated Frequencies in the Third Column Were Scaled with a Factor 0.8997; the Labeling of the Vibrational Modes Follows That of Bist28)
assignment
frequency/ cm-1 (calculated)
frequency/ cm-1 (scaled)
F1 τ β1 σ 11 F2 β2 16a 18b 16b 6a 6b 4 10b 12 torsion 10a 17b NH3.2 1 17a 18a 5 15 9b 9a 14 7a OH.bend 3 19b 19a 8b 8a NH3.4a NH3.4b 13.CH.st 7b.CH.st 2.CH.st 20b.CH.st 20a.CH.st OH.st NH3.1 NH3.3a NH3.3b
35 39 67 184 263 291 355 477 493 578 590 702 791 874 899 908 969 1043 1060 1109 1135 1136 1163 1194 1279 1311 1351 1402 1452 1542 1648 1687 1792 1816 1841 1843 3349 3359 3372 3386 3401 3700 3727 3896 3901
31 35 60 166 236 268 320 429 443 520
a
632 712 786 809 817 872 938 954 998 1021 1022 1046 1074 1150 1179 1215 1261 1307 1387 1483 1518 1612 1634 1656 1658 3013 3022 3034 3046 3060 3329 3353 3505 3510
IR intensity 5 3 5 6 4 3 25 0 13 4 1 1 14 97 20 257 18 27 629 6 0 1 2 7 18 7 79 133 34 94 61 29 50 47 43 50 4 9 44 29 8 823 99 22 20
frequency/ cm-1 (observed)
60,b 62c 164,b 162c
1.378 0.950 116.9 114.8
HF/6-31G
HF/6-31G(D,P)
HF/6-31G+(2D,P)
HF/6-31G+(2D,2P) 525b
822b
996,b 100.4d 1025.0d
1279.5d
3058a 3083a 3294a 3333a
TABLE 2: Structural Parameters of Phenol Clusters Calculated with the 6-31G Basis (Å and degrees) phenol
TABLE 3: Vibrational Frequencies, IR Intensities, and Their Assignments of Phenol-NH3 with Various Basis Sets
322b
This work. b Reference 22. c Reference 11. d Reference 7.
r1 r2 r3 R β γ
J. Phys. Chem., Vol. 100, No. 40, 1996 16055
phenol-H2O
phenol-NH3
1.367 0.960 1.819 117.4 115.8 178.9
1.364 0.996 1.914 117.5 116.4 171.7
results. In their calculations, much larger basis sets were used than those used in this work. Their simulation of high-frequency vibrations, however, did not fit very well with the spectrum observed in the present study. In particular, OH and NH vibrations appear at much higher frequencies than the observed, and even their frequency order was not appropriate. To confirm this discrepancy, we also have carried out the simulation of IR bands based on calculations with rather larger basis sets including polarization functions, that is, HF/6-31G(D,P), HF/
assignment
frequency/cm-1
intensity
OH NH NH NH NH NH NH OH NH NH NH OH NH NH NH OH
3700 3727 3896 3901 3704 3838 3840 3955 3693 3821 3823 3976 3696 3822 3824 3975
823 99 22 20 8 10 9 748 8 16 16 739 3 16 15 739
6-31G+(2D,P), and HF/6-31G+(2D,2P). For comparison the results of high-frequency vibrations are given in Table 3 together with other results. The calculated result, the same as that given by Schiefke et al.,with larger basis sets leads to a much higher frequency of the OH vibration than those of the NH vibrations and results in weaker intensities of the latter vibrations, being compared with the results obtained with smaller basis sets. Such a basis set dependence of calculated vibrational frequencies has been reported by Schu¨tz et al.8 for the phenol-H2O system. In this respect, therefore, an accidental nature may be involved in the agreement between the observed and the calculated frequencies based on HF/6-31G given in the present work. Despite the basis set effect, however, it is evident that all the calculated results lead to the largest intensity of the OH vibration. In this respect we have concluded that the observed IR band at 3294 cm-1 is associated with the OH stretching vibration of the phenol site and the weak band at 3333 cm-1 is due to one of the NH stretching vibrations. In the IR spectrum of phenol-NH3 shown in Figure 2a, it should be noted that the bandwidth of the OH band is substantially large, while that of the NH band is rather small. To avoid the saturation broadening of the transitions, the spectrum was measured under a sufficiently low fluence condition of the IR light. The full width at the half-maximum (fwhm) of the OH band at 3294 cm-1 is found to be 9 cm-1, while the fwhm of the NH band at 3333 cm-1 is less than 3 cm-1. The rotational temperature of jet-cooled phenol-NH3 was estimated to be less than 10 K from the line width of the 0-0 band of the S1-S0 transition. Thus, the rotational envelope of the OH band is not responsible for the line width. Such a large line width is commonly found for hydrogen-bonded OH stretching vibrations. The bandwidth of the hydrogen-bonded OH oscillator is larger than that of the NH oscillator, though the vibrational frequencies of these two bands are very close to each other. The CH bands show narrow widths. Therefore, it is suggested that the line broadening of the OH band is modeselective and does not depend on the density of states. This means that the characteristic relaxation process may be associated with the hydrogen-bonded OH oscillator. A theory for the mode-selective relaxation process will be discussed in a separate paper.26 2. Phenol-N(CH3)3. Figure 4 shows the IDIR spectra of (a) phenol-N(CH3)3, (b) phenol-NH(C2H5)2, (c) phenolN(C2H5)3, and (d) an isomer of phenol-N(C2H5)3. The spectra involving amines are similar to each other, but are quite different from that of phenol-NH3. A characteristic feature of the phenol-amine spectra is that intense bands appear in the 27003100 cm-1 region, and no band is seen above 3300 cm-1. Since
16056 J. Phys. Chem., Vol. 100, No. 40, 1996
Iwasaki et al.
Figure 4. IDIR spectra of (a) phenol-N(CH3)3, (b) phenol-NH(C2H5)2, (c) phenol-N(C2H5)3, and (d) an isomer of phenol-N(C2H5)3. Little circles in the spectra represent the OH stretch bands.
Figure 5. IDIR spectra of (a) phenol-h6-N(CH3)3 and (b) phenold5-N(CH3)3.
the PAs of these amines are much larger than that of NH3, it is reasonable to assume that phenol acts as a proton donor and amines are proton acceptors in these clusters. As seen in the IR spectrum of phenol-NH3, it is evident that the OH stretching frequency of the phenol site is largely red-shifted, while vibrational frequencies of the proton-accepting site are slightly affected by the hydrogen bond formation. The IR spectra of phenol-amines can be analyzed on the basis of these characteristics. It has been established from the IR absorption spectrum of N(CH3)3 in the gas phase that its symmetric CH stretching vibrations occur at around 2800 cm-1, and the antisymmetric vibrations appear in the region slightly below 3000 cm-1.27 The two distinct bands of phenol-N(CH3)3 at 2795 and 2844 cm-1 should correspond to the symmetric CH stretching vibrations of the N(CH3)3 site, and some bands near 3000 cm-1 should be attributed to the antisymmetric CH stretching vibrations. On the other hand, CH stretching vibrations of a phenyl ring would appear in the 3000-3100 cm-1 region, so that several bands above 3000 cm-1 in the phenol-N(CH3)3 spectrum are assigned to the CH vibrations of the phenol moiety. Despite the expectation that OH vibrations are the most intense in the IR spectra, there seems to be no remarkably intense band that can be attributed to the OH stretching vibration of the phenol moiety. Since the PA of N(CH3)3 (255.9 kcal/mol) is much larger than that of NH3 (204 kcal/mol),25 it is readily expected that the frequency of the OH band of phenol-N(CH3)3 is subjected to a larger red-shift than that of phenol-NH3. As a result, the OH stretching frequency becomes very close to those of the phenyl CH stretching vibrations and the strong intensity of the OH vibration may be distributed to the CH vibrations through the coupling among them. To confirm this, we have observed the IR spectrum of phenol-d5 (C6D5OH)-N(CH3)3, in which the CD vibrational frequencies of the phenyl ring are below 2400 cm-1, so that the congestion between the OH and CH vibrations is excluded from the observed region. Figure 5 shows a comparison of the spectra between phenol-d5-N(CH3)3 and phenol-h6-N(CH3)3. As seen in the spectrum of phenol-d5N(CH3)3, the bands above 3000 cm-1 disappear except for the broad and intense band at 3067 cm-1, while the CH bands of the amine moiety are mostly unchanged. Therefore, we assign the band at 3067 cm-1 to the OH vibration of the phenol moiety. Of course, in the case of phenol-h6-N(CH3)3, this vibration is no longer localized at the OH group and couples with the motion of CH groups. 3. Phenol-NH(C2H5)2 and Phenol-N(C2H5)3. The infrared spectra of phenol-NH(C2H5)2 and phenol-N(C2H5)3 are shown in Figure 4b-d. Similar to the spectrum of phenolN(CH3)3, several bands below 3000 cm-1 are assigned to the CH stretching vibrations of the amine moiety. In the phenol-
NH(C2H5)2 spectrum, it is difficult to give unambiguous assignments of OH and CH bands of the phenol moiety appearing in the 2900-3100 cm-1 region. Since the PAs are 225.1 and 225.9 kcal mol-1 for N(CH3)3 and NH(C2H5)2, respectively,25 the OH vibrational frequency of phenol-NH(C2H5)2 is expected to be very close to that of phenol-N(CH3)3. Thus, we tentatively assign the broad absorption at 3060 cm-1 to the OH vibration and two sharp bands at 3029 and 3089 cm-1 to the CH vibrations of the phenyl ring, though both of them may be coupled with each other. We also assign the bands at 2846 cm-1 to the symmetric CH stretching vibration and at 2974 and 2990 cm-1 to the antisymmetric vibrations of the amine moiety. In Figure 4c,d, two IR spectra are given for phenolN(C2H5)3. It has been known that there are at least two isomers of phenol-N(C2H5)3, as pointed out by Jouvet et al.13 They observed the S1-S0 electronic transition of phenol-N(C2H5)3 and found two intense bands at 34 460 and 35 631 cm-1, which have tentatively been assigned as 0-0 bands of isomers. In the present study, we tuned the probe UV laser wavelength to each 0-0 band. The two IR spectra obtained are shown in Figures 4c,d; the spectra c and d are obtained by tuning UV laser to the lower and higher 0-0 bands, respectively. The spectral features of the two spectra are similar to those of other phenol-amine clusters, and both spectra are similar to but not the same as each other, confirming the tentative assignment of the electronic spectrum given by Jouvet et al.13 Although the assignment of the OH stretching vibrations is not straightforward, we tentatively assigned them as broad and intense bands at 2985 and at 2990 cm-1 in Figure 4c,d, respectively. These bands are considered to be overlapped with the bands due to the antisymmetric CH stretching vibrations. The differences between these two spectra of the isomers are most notable in the CH stretching region around 2800 cm-1. This suggests that the structural difference between the isomers is due to the conformation of the ethyl groups. 4. The Red-Shifts of the OH Stretching Vibrations. The above assignments of the OH stretching vibration of phenolamine represent that extremely large red-shifts, which are more than 500 cm-1 from that of bare phenol, occur with the cluster formation. The red-shifts of phenol-amine clusters are much larger than that of phenol-NH3. Since amines are much more basic than NH3, this result suggests a correlation of red-shifts of hydrogen-bonded OH stretching frequencies with respect to PAs of accepting molecules. Such a correlation has been wellknown in vibrational spectroscopic studies in condensed phases. Recently, Jouvet et al.13 have presented a correlation between the 0-0 band shifts of the S1-S0 transitions and the proton affinities of the accepting molecules for the phenol-amine clusters. In Table 4 are listed the OH vibrational frequencies
IR Spectroscopy of Hydrogen-Bonded Phenol-Amine Clusters TABLE 4: Observed OH Vibrational Frequencies of Phenol Clusters, Their Shifts from Bare Phenol, 0-0 Band Frequencies of S1-S0 Electronic Transitions, Their Shifts from Bare Phenol, and Proton Affinities (PA) of Acceptor Molecules
phenol phenol-H2O phenol-CH3OH phenol-NH3 phenol-N(CH3)3 phenol-NH(C2H5)2 phenol-N(C2H5)3 isomer 1 phenol-N(C2H5)3 isomer 2
νOH/ cm-1
∆νOH/ cm-1
νel/ cm-1
∆νel/ PA/kcal cm-1 mol-1 g
3657a 3524a 3456b 3294 3067 ( 5 3060 ( 10 2985
-133 -201 -363 -590 ( 5 -597 ( 10 -672
36 348c 35 993c 35 933c 35 717d 35 532e 35 518f 35 460f
-355 -415 -631 -816 -830 -888
166.5 181.9 204 225.1 225.9 232.3
2990
-667
35 651f -697
232.3
a Reference 1. b Reference 4. c Reference 20. d Reference 22. e Reference 21. f Reference 13. g Reference 25.
J. Phys. Chem., Vol. 100, No. 40, 1996 16057 Conclusion In this work, the OH stretching vibrations of the phenolamine clusters were investigated by IR-UV double-resonance spectroscopy. We also performed the ab initio calculation on phenol-NH3, and the energy-optimized structure reproduced the observed spectrum very well. The OH frequencies of the clusters showed the extremely large red-shifts with increasing PAs of the amine molecules. We found a clear correlation between the red-shifts and the PAs of amines. It was also shown that the hydrogen-bonded OH stretching vibrations of the phenol-amine clusters strongly couple with the CH stretching vibrations. The substantial broadening of the hydrogen-bonded OH stretching bandwidths indicates a mode-selective vibrational relaxation process. The detailed and theoretical analysis of the band broadening is now in progress in our laboratory. References and Notes
Figure 6. Correlation between (a) red-shifts of the OH vibrational frequencies, (b) 0-0 band shifts of the S1-S0 electronic transitions and PAs of acceptor molecules.
(νOH) of the various hydrogen-bonded phenol clusters and their red-shifts (∆νOH) from νOH of bare phenol. For a comparison, Table 4 also lists the 0-0 band frequencies (νel) of the S1-S0 transition and their shifts (∆νel) from νel of bare phenol. Figure 6 shows the correlation of ∆νOH and ∆νel vs PA of accepting molecules. A similar correlation was reported by Hartland et al.,7 though their result was limited only to acceptors with rather small PAs. The present result shows that this correlation is valid over a much wider range of PA. It is concluded that the OH frequency shift is also an indicator of hydrogen bond strength, as well as the 0-0 band shifts of the electronic transition. When we know PAs of acceptor molecules, we can predict OH frequency shifts of clusters from this correlation. The excellent correlation between the red-shift and PA also indicates that it is not necessary to take into account intracluster proton transfer. If the intracluster proton transfer occurs, a hydrogen-bonded NH stretching vibration should appear instead of the OH stretching vibration, and the correlation should be destroyed because of the frequency difference between OH and NH stretching vibrations. Therefore, it is concluded that no proton transfer occurs in the ground state of these clusters involving amines with large PAs. This conclusion is consistent with the result of the electronic spectra measurements by Jouvet et al.,13 in which the phenol chromophore is still observed in these clusters.
(1) Tanabe, S.; Ebata, T.; Fujii, M.; Mikami, N. Chem. Phys. Lett. 1993, 215, 347. (2) Ebata, T.; Watanabe, T.; Mikami, N. J. Phys. Chem. 1995, 99, 5761. (3) Ebata, T.; Mizuochi, N.; Watanabe, T.; Mikami, N. J. Phys. Chem. 1996, 100, 546. (4) Watanabe, T.; Ebata, T.; Tanabe, S.; Mikami, N. J. Chem. Phys. 1996, 105, 408. (5) Pribble, R. N.; Zwier, T. S. Science 1994, 265, 75. (6) Pribble, R. N.; Garrett, A. W.; Haber, K.; Zwier, T. S. J. Chem. Phys. 1995, 103, 531. (7) Hartland, G. V.; Henson, B. F.; Venturo, V. A.; Felker, P. M. J. Phys. Chem. 1992, 96, 1164. (8) Schu¨tz, M.; Bu¨rgi, T.; Leutwyler, S. J. Mol. Struct. (THEOCHEM) 1992, 276, 117. (9) Bu¨rgi, T.; Schu¨tz, M.; Leutwyler, S. J. Chem. Phys. 1995, 103, 6350. (10) Gerhards, M.; Kleinermanns, K. J. Chem. Phys. 1995, 103, 7392. (11) Schiefke, A.; Deusen, C.; Jacoby, C.; Gerhards, M.; Schmitt, M.; Kleinermanns, K.; Hering, P. J. Chem. Phys. 1995, 102, 9197. (12) Solgadi, D.; Jouvet, C.; Tramer, A. J. Phys. Chem. 1988, 92, 3313. (13) Jouvet, C; Lardeux-Dedonder, C.; Richard-Viard, M.; Solgadi, D.; Tramer, A. J. Phys. Chem. 1990, 94, 5041. (14) Steadman, J.; Syage, J. A. J. Chem. Phys. 1990, 92, 4630. (15) Steadman, J.; Syage, J. A. J. Am. Chem. Soc. 1991, 113, 6786. (16) Steadman, J.; Syage, J. A. J. Phys. Chem. 1992, 96, 9606. (17) Syage, J. A. J. Phys. Chem. 1993, 97, 12523. (18) Hineman, M. F.; Kelley, D. F.; Bernstein, E. R. J. Chem. Phys. 1993, 99, 4533. (19) Kim, S. K.; Breen, J. J.; Willberg, D. M.; Peng, L. W.; Heikal, A.; Syage, J. A.; Zewail, A. H. J. Phys. Chem. 1995, 99, 7421, and references therein. (20) Abe, H.; Mikami, N.; Ito, M. J. Phys. Chem. 1982, 86, 1768. (21) Mikami, N.; Suzuki, I.; Okabe, A. J. Phys. Chem. 1987, 91, 5242. (22) Mikami, N.; Okabe, A.; Suzuki, I. J. Phys. Chem. 1988, 92, 1858. (23) Angstl, R.; Finsterho¨lzl, H.; Frunder, H.; Illig, D.; Papouek, D.; Pracna, P.; Narahari Rao, K.; Schro¨tter, H. W.; Urban, S. J. Mol. Spectrosc. 1985, 114, 454. (24) Frisch, M. J.; Trucks, G. W.; Head-Gordon, M.; Gill, P. M. W.; Wong, M. W.; Foresman, J. B.; Johnson, B. G.; Schlegel, H. B.; Robb, M. A.; Replogle, E. S.; Gomperts, R.; Andres, J. L.; Raghavachari, K.; Binkley, J. S.; Gonzalez, C.; Martin, R. L.; Fox, D. J.; Defrees, D. J.; Baker, J.; Stewart, J. J. P.; Pople, J. A. GAUSSIAN 92, Revision A; Gaussian, Inc.: Pittsburgh, PA, 1992. (25) Lias, S. G.; Liebman, J. F.; Levin, R. D. J. Phys. Chem. Ref. Data 1984, 13, 695. (26) Mikami, N. In preparation. (27) Bauer, S. H.; Blander, M. J. Mol. Spectrosc. 1950, 3, 132. (28) Bist, H, D,; Brand, J. C. D.; Williams, D. R. J. Mol. Spectrosc. 1966, 21, 76.
JP960711C
