Pressure Dependence of OH Stretching Vibrations. 2. Complexes of

methylated benzenes, and chloro- and bromobenzene shows a red shift over the ... higher than 4 GPa, whereas the low-frequency component is red shifted...
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J . Phys. Chem. 1991,95, 68-74

Pressure Dependence of OH Stretching Vibrations. 2. Complexes of Perfluoro-fert-butyl Alcohol with Aromatic Acceptors Thomas F. Mentelt and Werner A. P. Luck* Fachbereich Physikalische Chemie, Universitat Marburg, 0-3550MarburglLahn, FRG (Received: March 13, 1990; In Final Form: May 18, 1990)

The OH stretching vibration of "1:l" complexes of perfluoro-tert-butanol (PFTB) with I 1 aromatic acceptors dissolved in poly(ch1orotrifluoroethene) oil have been investigated as a function of pressure up to 8 GPa. Two types of frequency shift versus pressure functions have been observed: (1) With increasing pressure, the OH frequency of the complexes with benzene, methylated benzenes, and chloro- and bromobenzene shows a red shift over the entire pressure range. At about 8 GPa, the frequency shift is rougly double the frequency shift, Avap (Av = gas-phase frequency - frequency in the solution), at ambient pressure. (2)The OH frequency of the complexes with fluorinated benzenes shifts to the red by about 20 cm-l between 1 bar and 4 GPa and becomes pressure independent at higher pressures. The OH band of the system PFTB/fluorobenzene splits, with increasing pressure, into two components. The high-frequency component becomes pressure independent at pressures higher than 4 GPa, whereas the low-frequency component is red shifted at all pressures. The two types of observed Av versus pressure functions could be explained by specific interactions of the OH group to the fluorine atoms or to the r-system, respectively. The pressure vs Av curves for the systems in which the OH group interacts with the r-system fall on a single line when the relative frequency shift Av(p)/Avapis plotted versus the pressure. This indicates that the interaction potentials for these systems have the same form, differing only in the well depth, which is proportional to AvaP. Together with the results from part I I we find three classes of pressure behavior that are probably due to three different intermolecular potential forms. The short-range OH-acceptor interaction. which is of repulsive character in the van der Waals systems, is attractive in the case of the OHw-complexes.

Introduction Spectroscopy at high pressures has proven to be a useful tool for the study of intermolecular interactions in the condensed phase. The investigation of electronic transitionsZas a function of pressure provides, for example, information about the relative stabilization/destabilization of electronic states by intermolecular interactions, and the pressure dependence of vibrational transitions can reveal the perturbation of the vibrational potential of the electronic ground state by intermolecular interactions. Pioneering work in explaining the pressure dependence of X-H stretching vibrations in terms of the nature of intermolecular potentials was carried out by Drickamer and c o - w ~ r k e r s . ~ -The ~ results of Drickamer and co-workers for C-H and 0-H stretching vibrations have been rediscussed recently by Zakin and Herschbach6 from the viewpoint of the Schweizer-Chandler theory.' In a previous paper' (part 1 ) we described a spectroscopic method to investigate weak intermolecular interactions by highpressure infrared spectroscopy. We studied the pressure-induced OH-stretching frequency shifts, Av,of monomeric O H groups of 2,2,3,3,4,4,4-heptafluorobutyl alcohol (HFB) and perfluorotert-butyl alcohol (PFTB) in dilute solutions of the nonpolar solvents perfluoroheptane (PFHP) and poly(ch1orotrifluoroethene)(PCFE). We observed with increasing pressure ( I ) an increasing red shift at lower overall pressures, (2) a subsequent characteristic maximum shift Avmax,followed by (3)a quasilinear decreasing shift until the absorption maximum coincides with the gas-phase frequency ( A v = 0), and (4)a negative shift at the highest pressures (see the bottom of Figure 4). The pressureinduced maximum shifts Au,,, are proportional to the solventinduced frequency shifts Avap at ambient pressure, which are a measure of the interaction strength between solvent and the O H group.*-I0 These features of the pressure behavior of the O H stretching vibration could be described by a simple linear model." The model is based on the numerical evaluationI2 of the semiclassical phase integral $p, dg (p, momentum; g, normal coordinate) for the oscillation of the H atom resulting from the sum of the gas-phase O H potential and a perturbation potential, which accounts for the solvent effect. Because the interactions between the monomeric OH groups and the solvent environment were 'Present address: Department of Chemistry, Princeton University, Princeton. NJ 08544.

assumed to be of the van der Waals type, Lennard-Jones potentials were chosen to describe the solvent effect. Interaction potentials were determined1J3v14that give the correct Avmaxat an intermolecular distance R,, that is consistent with the distance one would expect from the Av vs R relationships in crystal hydrates. As one result, the model calculations show that A v ( R ) is linearly dependent on the gradient of the effective part of the intermolecular potential at R, i.e., on the intermolecular force. On the other hand, the calculated potential depths of the model O H solvent Lennard-Jones potentials (intermolecular energies) are linearly related to the experimental Av,p.1314The latter result is in agreement with the empirical findings, which establish linear relationships between frequency shifts and intermolecular interaction energies for van der Waals interactions8-l0 as well as for hydrogen bonded systems.15J6 However, the estimated effective potentials describe essentially the polarization of the OH group by a mean solvent environment; additional contributions to the total intermolecular energy, which do not affect the O H frequency, have not been considered.' The motivation for the high-pressure study presented here came 1312.17~18

( I ) Mentel, T. F.; Luck, W. A. P. J . Phys. Chem. 1990, 94, 1059. (2) Drickamer, H. G.In!. Reu. Phys. Chem. 1982,2, 171. Pressure Tuning Spectroscopy. In High Pressure Chemistry and Biochemistry; van Eldik, R., Jonas, J., Eds.; D. Reidel: Dordrecht, 1987; p 263. (3) Fishman, E.; Drickamer, H. G.J . Chem. Phys. 1956, 24, 548. (4) Benson, A. M.; Drickamer, H. G.Discuss. Faraday SOC.1956,22,39; J . Chem. Phys. 1957, 27, 1164. ( 5 ) Moon, S. H.; Drickamer, H. G.J . Chem. Phys. 1974, 61, 48. (6) Zakin, M. R.; Herschbach, D. R. J . Chem. Phys. 1988, 89, 2380. (7) Schweizer, K. S.; Chandler, D. J . Chem. Phys. 1982, 76, 2296. (8) Behrens-Griesenbach. A,; Schrems, 0.;Luck, W. A. P. J . Chem. SOC., Faraday Trans. 2 1984, 80, 579. (9) Luck, W. A. P.; Zheng, H. Y . J . Chem. SOC.,Faraday Trans. 2 1984, 80, 1253; Z . Naturforsch. 1984, 39A, 888. ( I O ) Kleeberg, H.; KoGak, 0.;Luck, W. A. P. J . S o h . Chem. 1982, 11, 611. ( 1 1) Mentel, T. F.; Peil, S.; Schioberg, D.; Luck, W. A. P. J . Mol. Struct. 1986, 143, 321. ( I 2) Schioberg, D. Habilitationsschrift; Philipps Universitat, Marburg, 1987. ( 1 3) Mentel, T. F. Dissertation, Philipps Universitat, Marburg, 1988. (14) Mentel, T. F.; Kiimmel, W. J . Mol. Liquids, in press. (15) Purcel, K. F.; Stikeleather, J. A,; Bunk, S. D. J . Am. Chem. SOC. 1969, 91,4019. Sherry, A. D.; Purcell, K. F. J . Phys. Chem. 1970, 74, 3535. (16) Sherry, A. D.; Purcell, K. F. J . Am. Chem. SOC.1972, 94, 1853. ( I 7) Lippincott, E. R.; Schroeder, R. J . Phys. Chem. 1955,23, 1099. Reid, C . J . Chem. Phys. 1959, IO, 182. (18) Saitoh, T.; Mori, K.; Itoh, R. Chem. Phys. 1981, 60, 161.

0022-365419112095-0068$02.50/0 0 1991 American Chemical Society

Pressure Dependence of OH Stretching Vibrations "" I

The Journal of Physical Chemistry, Vol. 95, No. 1. 1991 69 I

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Figure 1. Correlation between the frequency shift and the integrated intensity j c du of solutions of 2,2,3,3-tetrafluoropropanol in nonpolar solvents (left) and hydrogen-bond acceptors (right). The slope of the linear relation of the van der Waals systems is 2.0 X IO' cm2 mol-'; the slope of the linear relation for the hydrogen bond systems is 2.2 X IO5 cm2 mol-'. The solvents were perfluorohexane, 1,I ,2-trichlorotrifluoroethane, n-hexane, CCI4, dichloromethane, 1,l -dichloroethane, 1,2-dichloroethane, 1,2-dibromoethane, nitromethane, acetonitrile, and 1,4dioxane.

from two preceding results obtained by our group that pointed to a qualitative difference between van der Waals interactions and hydrogen bonds. First, the correlation between the energy of evaporation and the overtone O H frequency shift Auo2 of ethanol in solvents with different interaction strengthslO revealed two linear regimes, one for the van der Waals systems (dAuo2/dAUvap= 3 cm-' kJ-' mol, AuO2 < 70 cm-') and another for the H-bonded systems (dAuo2/dAU,,, = 80 cm-l kJ-l mol, AuO2 > 200 cm-I). For Auozbetween 70 and 200 cm-l the aromatic solvents benzene and bromobenzene form a kind of transition range between the two linear regions. Second, the study of the integrated intensity, I = J c du, as a function of the fundamental O H frequency shift for methanol highly diluted in different solvents, revealed two different linear regimes for small and large shifts.Ig Passing from the small Au of weakly interacting van der Waals systems to the larger Au of H-bonded systems, the magnitude of I and the slope dI/dAu show a sudden increase, of 50% and 20%, respectively, in the region between 60 and 70 cm-I. (A sudden increase of the half-width of the O H band was found in the same frequency range.I9) The two linear regimes are illustrated for the example of a fluoro alcohol, 2,2,3,3-tetrafluoropropanol,in Figure 1. In aromatic solvents benzene, toluene, mesitylene, and 1,2,3,5tetramethylbenzene, 2,2,3,3-tetrafluoropropanolhas frequency shifts between 75 and 1 IO cm-l and values of I between 15 X lo6 and 22 X IO6 cm mol-', which fall in between those of van der Waals and H-bonded systems shown in Figure 1. The slope dI/dAu in the case of the aromatic solvents is 2.5 X IO5 cm2 mol-'. Therefore we extended our high-pressure study of intermolecular interactions to complexes of alcohols with aromatic acceptors. To investigate this transition regime, we chose a series of halogenated and methylated benzenes as acceptor molecules. As a proton donor we used PFTB because of its low tendency to self-association,20 which reduces the competition between self-association and the desired 1 :1 complex formation. From gas-phase microwave results for H F and HCI 1 :1 complexes with benzene2, as well as from IR spectroscopic results for other alcohols interacting with aromatic acceptors in the condensed one can expect that (19) England-Kretzer, L.; Fritzsche, M.; Luck, W. A. P. J . Mol. Sfrucr. 1988, 175, 277. (20) Murto, J.; Kivinen, A.; Korppi-Tommola, J.; Viitala, R.; Hyomaki, J. Acta Chem. Scand. 1973, 27, 107. (21) Baiocchi, F. A,; Williams, J. H.; Klemperer, W. J . Phys. Chem. 1983, 87, 2079. Read, W.G.;Campbell, E. J.; Henderson, G. J. Chem. Phys. 1983, 78, 3501. (22) Szczepaniak, K.; Orville-Thomas, W. J. J . Chem. Soc., Faraday Trans. 2 1974, 70, 1175. (23) Pimentel, G.C.; McClellan, A. L. The Hydrogen Bond; Freeman: San Francisco, 1960; Chapter 6.4. (24) Andrews, L.; Johnson, G. L.; Davis, S.R. J . Phys. Chem. 1985,89, 1706.

TABLE I: Frequency Shift Av,, at Ambient Pressure and Composition of the Ternary Systems PRB/Aceeptor/PCFEa mole ratio vol ratio acceptor TFBZ 1,3-DFBZ 1,CDFBZ FBZ CBZ BBZ BZ TOL

AvJcm-' 52 59 68 80

MES

1I 4 I I4 137 147 158 170

TMBZ

178

XYL

PFTB/acceptor 1:7.5 1:7.5 1:8 1:8 1:8 1:8 1:7.5 1:7 1:8.5 1:6 1:2 1:5.5 1:ll 1 :4

acceptor/PCFE 1:4 1:4 1:4 1 :4 1:4 1:4 1:4 1 :4 1:4 1 :4 1:lO 1:4 1 :5 1 :5

TFBZ, 1,2,4-trifluorobenzene; 1,3-DFBZ, 1,3-difluorobenzene; IP-DFBZ, 1,4-difluorobenzene; FBZ, fluorobenzene; CBZ, chlorobenzene; BBZ, bromobenzene; BZ, benzene; TOL, toluene; XYL,pxylene; MES, mesitylene; TMBZ, 1,2,3,5-tetramethylbenzene.

in the case of PFTB the OH group will interact with the aromatic x-system. This opens up the possibility of changing the OH-xinteraction strength by varying the ring substituents. While the OH frequency shifts of the van der Waals systems in part 1 ranged from 8 to 30 cm-I, the series of substituted benzenes used in this study cover, for the AvaP of PFTB, the range 50-180 cm-I. Experimental Section The high-pressure IR spectra were recorded at about 40 OC in a modified diamond anvil cell13 of the NBS type.26 We used gaskets machined from hardened stainless steel foils (DIN 1.43IO) of 0.4-mm thickness. The cylindrical holes drilled through the gaskets to contain the solutions had a diameter of 0.3 mm. The high-pressure IR spectra were taken with a Perkin-Elmer 325 spectrophotometer equipped with an adjustable KBr beam condensore2' Pressures were measured by using the ruby fluorescence method.28 The pressure shift of the ruby R, signal was recorded with a Cary 82 Raman spectrometer using the 514.5-nm line of an argon ion laser (Coherent 52G) to excite the ruby fluorescence. With a single filling of the diamond anvil cell, a consecutive series of IR spectra was taken. Each spectrum was taken at an increased pressure. The new pressure was quickly checked before and finally determined after an IR spectrum. The scan speed was about 20 cm-'/min, which led to recording times of 20-30 min/IR spectrum. The reproducibility was checked in a few cases by rescanning the IR spectrum without changing the pressure. The whole procedure, including the pressure determination, of taking a single high-pressure spectrum required 40-50 min. The aromatic molecules chosen as acceptors were 1,2,4-trifluorobenzene (TFBZ), 1,3-difluorobenzene (1,3-DFBZ), 1,4difluorobenzene (1,4-DFBZ), fluorobenzene (FBZ), chlorobenzene (CBZ), bromobenzene (BBZ), benzene (BZ); toluene (TOL), 1,bxylene (XYL), mesitylene (MES), and 1,2,3$tetramethylbenzene (TMBZ). The invesigations were performed in ternary solutions consisting of PFTB, the acceptor, and PCFE. In most of the cases the mole ratio of PFTB to aromatic acceptor was chosen to be 1:7 to 1:8 and the volume ratio of acceptor to PCFE 1:4 or 1 5 . The composition of the mixtures are given in Table I together with the OH frequency shift Auap of the PFTB acceptor complexes in PCFE at ambient conditions. Dissolving the PFTB/acceptor mixture in PCFE prevented crystallization at high pressure, which occurs in the binary mixtures and cause losses of spectroscopic transparency and optical ho(25) Engdahl, A,; Nelander, B. J. Phys. Chem. 1985,89,2860; 1987,91, 2253. (26) Piermarini, G. J.; Block, S. Reu. Sei. Insfrum. 1975. 46, 973. (27) Adams, D. M.; Sharma, S . K. J . Phys. E 1977, 10, 838. (28) Piermarini, G. J.; Block, S.;Barnett, J. D.; Forman, R. A. J. Appl. Phys. 1915, 46, 2174.

70 The Journal of Physical Chemistry, Vol. 95, No. 1, 1991 mogeneity. The properties of PCFE as a pressure medium have been described earlier.'*29 At ambient conditions PCFE is a viscous liquid. It changes into a glassy state in the range 3-4 GPa and remains a transparent glass up to 12 GPa. The ternary solutions became very slightly opaque at pressures of about 3.5 GPa, which could be seen under a stereomicroscope, but stayed translucent up to the highest pressures. We interpreted this phenomenon as the same "glass transition" as found for the pure PCFE or the binary PFTB/PCFE solution.' For the system PFTB/xylene/PCFE we had to vary the PFTB to xylene mole ratio to I :2 and xylene to PCFE volume ratio to 1:10 because a mixture of 1:7.5/1:5 crystallized at 0.8 GPa. The 1:2/1:10 mixture showed the "glassy" state between 3.2 and 3.6 GPa. The disadvantage of the ternary mixtures is the competition between the OH-PCFE interaction and the OH-acceptor interaction. From the OH-PCFE interaction arises an IR band in the region of 3600 cm-I; however, with the exception of the system PFTB/TFBZ/PCFE, the features of the OH-PCFE and OHacceptor bands are well resolved. Moreover, the pressure dependence of the OH-PCFE IR absorption band is known from the previous study.' The choice of possible mixtures is restricted by the short path length of the light in the diamond anvil cell. Thus the composition ratios of the ternary mixtures are the necessary compromise between the demands of the high-pressure technique and the quality of the IR spectra. The frequency shift of the OH stretching vibration band in these high-pressure IR spectra can be determined with an accuracy of f 2 cm-l. Two independent high-pressure series of the system PFTB/MES/PCFE with different compositions and three independent high-pressure series of the system PFTB/FBZ/PCFE with similar compositions (see Table I) revealed that the frequency shift for a given pressure is reproducible within f 3 cm-I. This is a little more than the uncertainty of a single measurement, because some loss of the volatile PFTB during the filling procedure of the diamond anvil cell changes the composition of the solution and thus slightly the O H frequency of the PFTB/acceptor complex. The IR spectra of the alcoholic solutions at ambient pressure were measured at 40 'C in temperature-controlled quartz Infrasil cuvettes of 0.2-mm path length (Fa. Hellma, Miillheim). PFTB (SCM Chemicals, Gainsville, FL) and PCFE (Merck, Darmstadt, FRG) were used without further purification. The aromatic acceptors were kept over molecular sieves (3-4-A pore diameter) before use to reduce water traces. Especially in the case of the system PFTB/FBZ/PCFE water traces are dragged in unavoidably during the filling process of the diamond anvil cell at the open air. The presence of water gives rise to a small amount of a PFTB/H,O complex, with a PFTB-OH frequency of ca. 3200 cm-' at ambient conditions and water-OH absorptions at ca. 3700 cm-I.

Results Two examples of series of IR spectra as a function of pressure are shown in Figures 2 and 3. In the system PFTB/BZ/PCFE (Figure 2), the O H stretching band of the PFTB/BZ complex at ~ 3 5 0 0cm-' is well separated from the band at ~ 3 6 0 0cm-' resulting from the OH-PCFE interaction. The maximum of the OH/benzene band shifts with increasing pressure continuously to lower wavenumbers. (The OH/PCFE band shows the same pressure dependence as in the binary system (shown in the lower part of Figure 4), which has been discussed in part I.') The half-width of the OH/benzene band increases with increasing pressure; the integral intensity shows a maximum at pressures of about 2 GPa. In the IR spectra of the system PFTB/I,CDFBZ/PCFE, which are shown as a function of the pressure in Figure 3, the O H band due to the PFTB/ 1 ,4-DFBZ interaction at ~ 3 5 5 cm-l 0 is partly overlapped by the O H band due to the PFTB/PCFE interaction ( ~ 3 6 0 0cm-I). Nevertheless the maximum of the PFTB/1,4(29) Langer, K; Luck, W. A. P.; Schrems, 0. Appl. Spectrosc. 1979.33,

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Figure 2. OH stretching vibrations of the system PFTB/benzene/PCFE as a function of the pressure. The absorption band of the PFTB/benzene complex occurs in the range 3500-3350 cm-I. The smaller feature at 3600 cm-l is due to the interaction of OH groups with PCFE. The pressure (in GPa) increases from the bottom to the top: 0.05,0.11, 0.46, 0.80, 1.31, 1.95, 2.54, 3.21, 3.68, 4.14, 4.81, 5.67, 6.56, 7.32, and 8.29. The frequency shift at ambient pressure is marked by the bar at the bottom. For the sake of clarity the IR spectra are stacked in this figure and in Figures 3 and 7. 2632

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Figure 3. OH stretching vibrations of the system PFTB/l,4-difluorobenzene/PCFE as a function of the pressure. At the absorption band of the PFTB/I,4-difluoro complex at 23550 cm-' overlaps with the absorption band at 3600 cm-I for PFTB interacting with PCFE. The pressure in GPa increases from the bottom to the top: 0.29, 0.55, 1.06, 1.76, 2.52, 2.77, 3.14, 3.63, 4.17, 4.74, 5.67, 6.37, 7.37, and 8.07. The frequency shift at ambient pressure is marked by the bar at the bottom.

DFBZ band is well recognizable. From ambient pressure to about 4 GPa this maximum is shifted to smaller wavenumbers. At

pressures higher than 2 4 GPa the frequency shift is approximately independent of the pressure. The effect of the pressure on the intensity and the half-width Avllz is qualitatively comparable to that of the PFTB/BZ/PCFE system, but the maximum of the

Pressure Dependence of O H Stretching Vibrations 350

The Journal of Physical Chemistry, Vol. 95, No. 1. 1991 71

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FBZ (OH-n) FBZ (OH-F) 1,4-DFBZ 1,3-DFBZ TFBZ PFTBIPCFE HFWPCFE PFTBIPFHP HFWPFHP 9 . .

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Figure 4. Frequency shift as a function of the pressure for ternary PFTB/acceptor/PCFE systems (group I , top, and group 2, middle) and binary fluoroalkyl alcohol/nonpolar solvent van der Waals systems (group 3, bottom).

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Figure 5. Pressure broadening of the OH stretching infrared bands of the OH/acceptor complex for the systems PFTB/ 1,4-DFBZ/PCFEand PFTB/TMBZ/PCFE. Note that the half-width A U * ~ as , ~ measured , at the low-wavenumber side, is given and not the full width.

band area (see below) which is again at about 2 GPa, is less pronounced. The pressure dependence of the O H frequency shifts Au of the complexes of PFTB with benzene and the substituted benzenes is shown i n the upper and the middle part of Figure 4. Considering the frequency shift as a function of the pressure, the aromatic acceptor molecules appear to form two groups: (1) benzene, the methylated benzenes, and chloro- and bromobenzene (Auap > 100 cm-I), which show large pressure-induced red shifts up to the highest investigated pressures ( 4GPa); (2) the fluorobenzenes, with ambient pressure shifts AvaP between 50 and 80 cm-I, for which the frequency shifts after a pressure-induced red shift of about 20 cm-I are approximately pressure independent at pressures larger than 4 GPa. In Figure 5 the increase of the half-width of the OH/acceptor band as a function of the pressure is demonstrated for PFTB/ TMBZ/PCFE as an example of group 1 systems and for PFTB/1,4-DFBZ/PCFE as an example of group 2 systems. As

1

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Figure 6. Pressure dependence of the integral intensity of the OH stretching infrared bands of the OH/acceptor complex for the systems PFTB/I ,CDFBZ/PCFE and PFTB/TMBZ/PCFE. As a measure for the band area the product of Av*,,* (roughly half of the full-width at half-maximum; in cm-I) and the band height (maximum absorbance; dimensionless) is used.

a measure of the pressure broadening, we chose the distance Au* from the band maximum position to the low-wavenumber s i l i at half-height. This side of the OH/acceptor band is, in the case of 1,4-DFBZ, less perturbed by the overlap with the PFTB/PCFE band. For better comparison we gave A U * ~for , ~ the system PFTB/TMBZ/PCFE, too. In the latter case A U * ~is, / ~within 1 cm-', half of the full width at half-height. The increase of the half-width of the O H stretching with increasing pressure is probably due to multiple influences. The loss of hydrostatic pressure conditions in the solutions that become glassy at p > 3.5 GPa is the probable cause of an inhomogeneous band broadening at higher pressures. One might speculate that a contribution of a homogeneous nature could arise from the shortening of the vibrational relaxation time caused by the increase of the strength of the intermolecular interaction with increasing pressure.30 A decrease of the vibrational relaxation with increasing intermolecular forces has been found by Heilweil et aL3I for SiOH interacting with different solvents and by Baggen et al.32for the N2 vibration under the influence of the attractive part of the intermolecular potential. Another possible contribution to a homogeneous broadening might be due to the increase of the steepness of the effective part of the intermolecular potential, so that a broader range of intermolecular forces is covered by the amplitude of the H m o t i ~ n . ~ ~ ~ ~ In Figure 6 the change of the integral intensity for the OH/ acceptor band is shown as a function of the pressure for the systems PFTB/ 1,4-DFBZ/PCFE and PFTB/TMBZ/PCFE. As a measure of the band area we chose the product of A V * ~and / ~ the height of the band. The band areas have a maximum at about 2 GPa. This result is more distinctly marked in the case of TMBZ. All other systems under consideration show qualitatively similar behavior with maxima of the band area at about 2 GPa. The change in the band areas with pressure is probably due to a combination of effects. Two factors that would produce an increase of the band area at lower pressures could be the intrinsic increase of the absorption coefficient with increasing red shift,I9 similar to the situation shown in Figure 1, and the increase of the number of OH-benzene complexes due to the preference of stronger interactions at higher pressure^.^ The breakage of OH-acceptor complexes by shear stresses in the "glassy" state could be a contribution to the decrease of the band area at higher pressures. Another contribution to the pressure dependence of the band areas originates in the highpressure technique. The pressure on the sample is generated by (30) Zerda, T. W.; Zerda, J . J . Phys. Chem. 1983, 87, 149. (31) Heilweil, E. J.; Casassa, M. P., Cavanagh, R. R.; Stephenson, J. C. J . Chem. Phys. 1985, 82, 5216. (32) Baggen, M.; van Exter, M.; Lagendijk, A. J . Chem. Phys. 1987.86, 2423. Baggen, M. Dissertation, Universiteit van Amsterdam, 1989, pp 84-87. (33) Buanam-Om, C.; Luck, W. A. P., Schioberg, D. Z.Phys. Chem. 1979, 117, 19.

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The Journal of Physical Chemistry, Vol. 95, No. I , 1991 -1.0

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3300 v / om-'

Figure 7. OH stretching vibrations of the system PFTB/fluorobenzene/PCFE as a function of pressure. The absorption band of the PFTB/fluorobenzene complex at ~ 3 5 5 0cm-l splits with increasing pressure. The low-frequency component is strongly red shifted with increasing pressure. The ambient pressure frequencies of both components from a band fit are given as bars at the bottom. The band at about 3600 cm-' is due to PFTB, which interacts with PCFE only. The small band at about 3700 cm-' originates from water traces. The pressure (in GPa) increases from the bottom to the top: 0.15,0.29, 0.41, 0.83. 1.26, 1.62, 2.02, 2.31, 2.71, 2.91, 3.31, 3.69, 4.60, 5.47, and 6.28.

squeezing the gasket between the diamond anvils. As a result, the cylindrical hole in the gasket that contains the sample is deformed. From the Lambert-Beer law A = tcl, wherein A is the absorbance, t is the absorption coefficient, 1 the length of the light path in the sample, and c the concentration, follows with c = n/Vand V = Vcylindric = i ~ ( d / 2 ) that ~ L A = t n ~ / ( d / 2 ) ~Thus, . for a given cylindrical sample, A is independent of the compression of the cylinder as long as the diameter remains constant. This is approximately true at lower pressures. The shortening of the light path I due to the compression of the gasket is then nearly compensated for by the increase of the density of the solution, which is equivalent with the increase of the concentration c. At higher pressures the metal gasket exhibits cold flow. This causes a distinct but continuous increase of the diameter d of the cylindrical hole in the gasket that contains the sample. The decrease of the path length and the simultaneous increase of the sample diameter cause a decrease of the absorbance A and therefore the decrease of the band area. Both the pressure dependence of the half-width and the band area will be mentioned here only as a result (with some possible explanations of the effects). The discussion will be concentrated on the frequency shift as a function of the pressure. The complex of PFTB with fluorobenzene is a special case. The OH stretching band splits in two subbands with increasing pressure (Figure 7). After a pressure-induced red shift of =25 cm-I, the main subband at ~ 3 5 5 cm-' 0 becomes pressure independent at p > 4 GPa, and the pressure dependence of its frequency shift fits well into the series of the trifluoro- and difluorobenzenes (see the middle of Figure 4). The pressure-induced red shift of the minor subband at ~ 3 4 5 cm-' 0 agrees with the Av =f(p) curves of the complexes of PFTB with benzene, chloro- and bromobenzene, and the methylbenzenes, defined as group 1 in the upper part of Figure 4. For the complexes of PFTB with the other monohalobenzenes, chloro- and bromobenzene, dissolved in PCFE, no obvious pressure-induced splitting was observed. However, the half-width of the O H stretching band of the bromobenzene complex is exceptionally large. The band at ~ 3 6 0 cm-l 0 in Figure

Mente1 and Luck 7 is again due to the PFTB/PCFE interaction, whereas the small band at ~ 3 7 0 cm-l 0 originates from water traces, as mentioned in the Experimental Section. The high-frequency wing of a PFTB/water band (with maximum at ca. 3200 cm-]) can be seen at the lowest wavenumbers in Figure 7 . This band is red shifted with increasing pressures, but a part of the high-frequency wing can be recognized at the low-wavenumber edge of all spectra in Figure 7. The combination of the results for aromatic acceptors and those for weakly interacting systems from part 1 ' yields the three different types of Av vs pressure functions as shown in Figure 4. This suggests three classes of intermolecular interactions: group I , with Av, > 100 cm-' characterized by continuous pressureinduced re( shifts, group 2 with Av,P(s between 40 and 80 cm-' characterized by the pressure independence of Av at p > 4 GPa, and group 3 of weak van der Waals interactions with AvaP< =30 cm-I, revealing a maximum AvmaXin frequency shift and a blue shift at high pressures.

Discussion The splitting of the OH stretching vibration of the PFTB/ fluorobenzene complex with increasing pressure allows us to explain why the aromatic acceptors under investigation show two types of pressure behavior. The OH frequency of the acidic alcohol PFTB is quite sensitive to differences in the solvent environment. A splitting of the OH vibration band of PFTB caused by specific interaction of the OH group with different parts of the solvent molecules has been found previously.34 An example of pressure-induced splitting has been described for the binary system PFTB/PCFE in part 1' (and see the OH-PCFE band at ~ 3 6 0 0 cm-' in the Figures 2, 3, and 7). A splitting of the OH stretching vibration band caused by two specific interactions has been found by Szczepaniak and Orville-Thomas22in a binary solution of phenol in fluorobenzene. They assigned the subband at higher wavenumbers (Avap = 50 cm-I) to an interaction of the OH group with the fluorine atom of fluorobenzene and the subband at lower wavenumbers (AvaP= 87 cm-I) to an interaction of the OH group with the aromatic mystem. The ability of substituted aromatic molecules to interact with a given donor through either the ring substituents or the aromatic mystem has been also found, e.g., the work of Wayland and drag^^^ on anisole and thioanisole. In our case a splitting of the OH band at ambient pressure is not obvious in the ternary system PFTB/FBZ/PCFE, but in the binary solution of PFTB in fluorobenzene we found a distinct shoulder at Avap = 137 cm-I besides the main band at Avap = 87 cm-'. Our Gaussian-Lorentzian sum-band analysisI3 of the OH stretching band in the ternary system PFTB/FBZ/PCFE yields a main band with AvaP = 79.5 f 2 cm-I and a small band with AvaP = 132 f 5 cm-I whose intensity is about one-tenth of the main band. The small differences in the frequency shifts of the binary and the ternary system are due to the different solvent environments of the PFTB/FBZ complex in the binary and the ternary solution, whereas the larger OH frequency shifts of PFTB with respect to those of phenol can be explained by the larger acidity of PFTB (pKa = 5.436) with respect to phenol (pK, = 9.9737). Therefore we assigned the high-frequency main component of the OH stretching band in the system PFTB/FBZ/ PCFE to an OH-F interaction and the low-frequency part to an interaction of the O H group with the mystem of fluorobenzene, which we will abbreviate as OHwr interaction. The assignment of the low-frequency subband to the OH--P interaction is strongly supported by the fact that the pressure dependence of its frequency shift agrees well with that of the complexes of PFTB with benzene and the methylbenzenes, in which an OH-T interaction is dominant (see upper part of Figure (34) Schriver, L.; Burneau. A . J. J . Chim. Phys. 1976, 73, 723; J . Chem. SOC.1985, 81, 503.

(35) Wayland, B. B.; Drago, R. S . J . Am. Chem. SOC.1964, 86, 5240. (36) Fuller, R.; Shure, R. M. J . Org. Chem. 1967, 32, 1217. (37) Murto, J. Hydroxide-Alkoxide Ion Equilibria and Their Influenceon Chemistry of Functional Groups. In The Chemistry of Funcfional Groups; Patai, S . , Ed.; Interscience: London, 1971; Vol. IO, Part 2, p 1087.

The Journal of Physical Chemistry, Vol. 95, No. 1, 1991 13

Pressure Dependence of O H Stretching Vibrations

I

j w

“

1

0

_II

. ,

VI

1.75

j I I

! 50

100

150

200

frequency shift I cm-1

Figure 8. Slope of the Au versus pressure functions at ambient pressure (initial slope) as a measure of the “rate of red shift” as a function of the frequency shift at ambient conditions. The initial slope (open triangle) for the OH-CI interaction’ in the system PFTB/PCFE has been added to the data of group 2. The large error bars for this system and for PFTB/TFBZ (open square) indicate that these results are from a band fit.

2.00

’

0

1

~

’

2

-

’

3

~

4

~

5

”

6

-

’

7

-

8

’

~

9

10

’

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pressure I GPa

Figure 9. Relative frequency shifts Av/Auap of the complexes of PFTB with benzene, toluene, xylene, mesitylene, 1,2,3,5-tetramethylbenzene, and fluorobenzene (OH-n) dissolved in PCFE as a function of pressure.

gradient of the interaction potential (a force) at that distance). For the acceptors benzene, the methylbenzenes, and fluorobenzene ( O H w r subband) the bunch of diverging curves of the group 1 in Figure 4 can be reduced to a single function by plotting the relative frequency shift (Au(p)/Auap) vs the pressure (Figure 9). The scatter of the data points in Figure 9 is about twice the uncertainty of the single measurements. Provided that for the OH-r systems a linear Avap-potential well depth relation exists, such as in the van der Waals case, the result shown in Figure 9 suggests that all these systems have the same form of the OH-r interaction potential and that the potentials differ only in the well depth. Or more precisely, if the interaction potentials, which determine the frequency shifts, have the usual form of a product of the well depth Dotimes a termf(R - Ro), and Av(R) is given by the gradient of the interaction potential at R [Le., Av(R) = Daf‘(R - Ro)], all systems under consideration would have the samef(R - Ro). Under these conditions it is not necessary that the minimum distances Ro of the interaction potentials are the same. However, one should keep in mind that the reduced relation in Figure 9 is plotted as a function of pressure and not as a function of intermolecular distance. So the statement above is true only if the pressure changes the intermolecular distance in the same way for all these systems, Le., if the systems have the same local compressibility in the region of the OH-.r interaction. This is not expected a priori for systems with different interaction strengths, although in general the compressibilities for most of the organic liquids are similar (to within ca. 15%). But in turn because the reduced Av/AvaP vs pressure curves shown in Figure 9 fall on a single line and the involved acceptors are chemically similar, one can assume that the local compressibilities are the “same”. Then the hypothesis would hold that the intermolecular potentials of these PFTB/acceptor complexes have the same analytical form, differing only in the well depth. And the frequency shift at ambient pressure would in fact be a linear measure of the well depth of the intermolecular potential, as in the case of the van der Waals systems. A linear correlation between Ava and the potential well depth would be consistent with the empirical linear relation between frequency shift and hydrogen bond enthalpy for complexes of PFTB with hydrogen-bond acceptors.17 The acceptors chloro- and bromobenzene do not fit on the common line of the other OH-r systems but seem to form a relation of their own. At 7 GPa the relative frequency shifts (Au(p)/Auap) for PFTB/CBZ and PFTB/BBZ are ca. 15% larger than the Av@)/Avap of the other PFTB complexes of group 1. This finding is confirmed by the slight deviation of chlorobenzene and bromobenzene from the common line for the OH-r systems in Figure 8. It has been found previously that these halobenzenes form an exception with respect to other aromatic acceptors.38 This

4). On the other hand, the pressure dependence of the frequency shift for the OH-F component of the PFTB/FBZ complex is similar to the Av vs pressure function of the PFTB complexes with 1,2,4-trifluorobenzene, 1,3-difluorobenzene, and 1,4-difluorobenzene (see the middle of Figure 4). This fact in turn suggests that in the cases of the difluoro- and trifluorobenzenes as well the OH group interacts with the fluorine atoms instead of the rsystem. A possible reason for a preference of the OH-F over the OH-r interaction in the complexes of PFTB with the fluorinated benzenes may be the reduction of the electron density of the aromatic system relative to benzene due to the electronwithdrawing effect of the fluorine atoms and the appearance of an appreciable dipolar components in the direction of the C-F bonds. Using this interpretation of our high-pressure results, we are able to assign the different types of Au vs p curves in Figure 4 to three different classes of acceptors: (1) acceptors with rsystems, which form hydrogen-bond-like OH-T interactions; (2) acceptors with fluorine atoms bound to a polarizable molecular residual such as the fluorobenzenes, which form OH-F interactions; (3) aliphatic acceptors with F atoms, which are bound to a “weak polarizable” molecular residual like the highly fluorinated alkanes. This classification is confirmed if the slope (dAu/dp),, of the Av versus pressure functions of Figure 4 at ambient pressure is plotted as a function of Avap, as shown in Figure 8. The van der Waals systems of group 3 fall roughly on one line, and the OHwr-systems on another line. The two lines have different slopes. The complexes PFTB/CBZ and PFTB/ BBZ fall slightly off the OH-r line (see also below). To the data of the PFTB complexes with the fluorobenzenes in Figure 8 (middle) we added (dAu/dp)apof the Av versus pressure function of the OH-.CI interaction in the system PFTB/PCFE (see part I I ) . The frequency shift of this specific interaction showed a pressure independence at high pressures, too. With the exception of the PFTB/TFBZ complex, the initial slopes of the group 2 systems seem to be independent of the frequency shift at ambient pressure. The overall tendency to a pressure-induced red shift increases from the van der Waals systems to the OH-r systems with increasing Auap. For the binary systems HFB/PFHP, PFTB/PFHP, HFB/ PCFE, and PFTB/PCFE (group 3) we estimated13J4interaction potentials, which give the correct features of the frequency shift as a function of the pressure. In this sequence of van der Waals systems the potential well depths increase from 220 to 840 cm-’, whereas the potential minimum distances decrease from 2.5 to 2.1 ?,.I Every system has its own interaction potential. However, we found a linear relation between the potential well depth (an energy) and the frequency shift AvaPat ambient c ~ n d i t i o n s ~ ~ ’ ~ J ~ (38) Josien, M.-L.; Sourisseau, G.Contribution A L‘Etude Des Molecules (besides the result from our model that the frequency shift for Aromatiques ConsidEries Comme Accepteurs De Protons. In Hydrogen Bonding; Hadzi, D., Ed.; Pergamon Press: London, 1959; p 129. a certain intermolecular distance R is linearly related to the

’

~

Mente1 and Luck

74 The Journal of Physical Chemistry, Vol. 95, No. 1 , 1991 could reflect a different potential form or different pressure intermolecular distance function. This in turn may be due to the bulkiness of the halogen substituents or a direct involvement of the chlorine and bromine atoms in the intermolecular interaction. The Au versus pressure curves of group 2 of the fluorobenzenes, including the high-frequency subband of fluorobenzene, cannot be reduced to a single function in the Av(p)/AvaP vs pressure plot. This is also true for the van der Waals systems of group 3, even though their potential well depths scale linearly with A U , ~ Let us assume that exerting pressure on a liquid sample simply reduces the mean intermolecular distances R . At low pressures, the pressure-induced red shifts of all systems (Figure 4) indicate that the attractive forces increase with decreasing intermolecular distance. The rate of red shift (Figure 8) is different for each system and roughly correlated with the interaction strength between OH group and acceptor molecule. In this pressure range, the pressure increase primarily compresses the liquid’s “free” volume, which is due to the temperature oscillation of the mole c u l e ~ . At ~ ~ higher pressure, say 5 GPa or more, we can reasonably assume that the mean intermolecular distance is equal in magnitude to the sum of the van der Waals radii of adjacent molecules (atoms) or smaller. The compression must then be performed against the total repulsive molecular forces. At these high pressures, short-range interactions are dominant. For the van der Waals systems of group 3, our model calculations show that the negative frequency shifts at high pressures are due to the effect of the repulsive branch of the intermolecular potential, in agreement with earlier results of Drickamer and c o - w o r k e r ~or~B~ckingham.~’ ~~~ This clearly confirms that the short-range O H solvent interactions in the van der Waals systems are of repulsive character. The influence of the repulsive forces starts around Aumax,and all of the attractive forces are compensated for when at high-pressure conditions the frequency of the gas phase occurs. In H-bonded OH-B systems in crystals, the H-B distance is smaller than the sum of the van der Waals radii of the H and the B atom even at normal pressure.42 In medium-strong H bonds ( A u > 300 cm-I), the 0.-B distance is smaller than the sum of the 0 and B van der Waals radii even though the H atom is located between them.42%43When one considers the OH-*-interaction as a form of hydrogen bonding,23it is apparent that these arguments apply to our systems as well. The frequency shifts of the OH-.lr-systems are positive (red shift) and increase with decreasing intermolecular distance at the highest densities for pressures of about 8 GPa. Therefore, contrary to the van der Waals systems, the local short-range interactions between the O H group and the acceptor molecule must be attractive in nature, though an increasing long-range attraction of the second-neighbor solvent shell, which also moves closer to the OH group at higher (39) Jorgensen, W. L.; Ibrahim, M. J . Am. Chem. Soc. 1982, 104, 373. (40) Wiederkehr. R. R.; Drickamer, H. G . J . Chem. Phys. 1958, 28, 31 I . (41) Buckingham, A. D. Proc. R. SOC.London 1958,248A, 169; J. Chem. Soc., Faraday Trans. 1960, 56, 753. (42) Mikenda, W. J . Mol. Struct. 1986, 147, I . (43) Olovsson, 1.; Jonsson, P . G . X-ray and Neutron Diffraction Studies of Hydrogen Bonded Systems. In The Hydrogen Bond; Schuster, P., Zundel, G . . Sandorfy, C.. Eds.; North Holland: Amsterdam, 1976; Vol. 2, p 393.

pressures, may contribute slightly to the total attractive interaction. The application of results of ab initio energy partitioning calculations for hydrogen-bonded system^^^-^^ to the OH--x systems suggests that the strong short-range attractive interaction could be of the “charge-transfer” type. Assuming that to a first approximation the frequency shift is a measure of the gradient of the interaction p ~ t e n t i a I , ’ ~ the !’~*~~ constancy of 4 v ( p ) for the complexes of PFTB with the fluorinated benzenes of group 2 seems to indicate that the potential is nearly linear in the range of intermolecular distances that are realized at pressures between 4 and 8 GPa. The pressure independence of the frequency shift could also be due simply to a very flat maximum, similar to the 4 u vs p curves for the van der Waals cases. To clarify this point, measurements at much higher pressure are necessary. In comparison with the van der Waals systems on one hand and the O H v r systems on the other, repulsive and attractive short-range forces in the OH-F-@ systems seem to be balanced. Conclusions

High-pressure infrared spectroscopy using the stretching frequency of the O H group as probe to monitor intermolecular interactions yields, in the case of the complexes of PFTB with aromatic acceptors, useful, if qualitative, results. The method is useful for detecting specific intermolecular interactions and thereby illuminating the microscopic structure of the ternary solutions PFTB/acceptor/PCFE. We were able to demonstrate that the character of the short-range intermolecular interaction is repulsive for the van der Waals systems and attractive for the H-bond-like OH.-* systems. A semiquantitative treatment of the complexes of PFTB with the aromatic acceptors similar to that carried out for the van der Waals systems turned out to be difficult, because of the absence of characteristic quantities, such as Aumax, that are independent of the pressure-intermolecular distance relation and to which potential parameters could be fit. A knowledge of the pressure versus intermolecular distance function, however, would enable us to carry out semiquantitative calculations, e.g., using a developed form of the LippincottSchroeder model.’* The lack of such knowledge forces us to be somewhat descriptive in our analysis, but we have been able to classify the types of interactions and intermolecular potentials for the present systems. Acknowledgment. We gratefully acknowledge the financial support of the high-pressure project by the Deutsche Forschungsgesellschaft. (44) Bratoz, S. Electronic Theories of Hydrogen Bonding. In Aduances in Quantum Chemistry; Lowdin, P. O., Ed.; Academic Press: New York, 1967; Vol. 3, p 209. (45) Schuster, P. Energy Surfaces for Hydrogen Bonded Systems. In The Hydrogen Bond; Schuster, P., Zundel, G.,Sandorfy, C., Eds.; North Holland: Amsterdam, 1976; Vol. 1, p 25. (46) Morokuma, K. Acc. Chem. Res. 1977, 10, 295. Morokuma, K.; Kitaura. K. Energy Decommsition Analysis of Molecular Interactions. In Chemical Appliciiions of Aiomic and M&cular Electrostatical Potentials; Pollitzer, P., Truhlar, D. G . , Eds.; Plenum Press: New York, 1981; p 215. (47) Liu, S.; Dykstra, C. E. J . Phys. Chem. 1986, 90, 3097.