J. Phys. Chem. 1983, 87,3664-3670
3664
can be said for the w3 assignment even if the g 3 value is comparable with that of g6 and g,. However, an assignment of w6 has been already ruled out and that of the ogmode cannot be compared, given its low ( w ” ~ = 337 cm-’) frequency value, with the experimental vibronic frequencies. The LiTCNQ PR5,6Raman spectra support our assignment of the w3 and o7vibrations since they are reported as the most strongly intensity enhanced. The calculated coupling constants allow one to reject the attribution of the vibronic structure of the visible absorption system to the 2AU-2B2g transition. In fact, the values reported in the fifth column of Table V suggest that the w g and w4 vibrational modes are the most heavily involved in a possible vibronic pattern of this transition. The first has been already excluded and the association of the w4 mode with the u, experimental frequency is improbable. In fact, this mode being practically insensitive to deuterium or fluorine substitution, this association would require a too large variation in the normal coordinates in going to the excited state. In conclusion, the transition which displays the vibronic structure in the visible spectral range corresponds to the
2B3u-2Bzg, coupled to the w3 and w7 vibrational modes. The fit between experimental assignment and calculated coupling constants is reasonably successful if one considers that the closeness of the two electronic transitions certainly reduces the reliability of the calculated value^.'^^^^ One can get more precise information through an absorption profile calculation, taking into account both the diagonal and the off-diagonal coupling constants, as well as the i~ bond order variations. Such calculation is planned for future work in this laboratory.
Acknowledgment. Financial support by the National Research Council and by the Ministry of the Education of Italy is acknowledged. We thank Drs. R. Bozio and A. Girlando for useful discussions. The technical assistance of Dr. C. Ricotta in the chemical syntheses has been much appreciated. Registry No. TCNQ radical anion, 34507-61-4; LiTCNQ-d4, 86392-67-8; LiTCNQ-F4,34473-33-1. (25) L. S. Cederbaum and W. Domcke, Chem. Phys. Lett., 25, 357 (1974).
Raman Spectroscopic Study of Aqueous LiX and CaX, Solutions (X = CI, Br, and I ) in the Glassy State H. Kanno” Department of Chemistty, Meisei Universiw, Hino, Tokyo 19 1, Japan
and J. Hlralshl National Chemical Laboratory for Industry, 1- 1 Yatabe, Tsukubagun, Ibarakl305, Japan (Received: October 27, 1982; In Final Form: March 16, 1983)
Raman spectra of aqueous LiX and CaX2 solutions (X = C1, Br, and I) in the glassy state were measured with special attention given to the low-frequencyRaman bands and to the Raman band intensity changes with anion or salt concentration. The importance of the charge-transfer states of hydrogen bonds OH...X- (X- = halide ion) is confirmed for the intensity changes of all the Raman bands in these halide solutions. It is shown that detection and characterization of low-frequencyRaman bands for aqueous solution are easily achieved in a Raman spectrum for the solution in the glassy state. Spectral changes associated with going from the liquid state to the glassy state were also studied for the OH stretching Raman spectra of the LiCl solutions.
Introduction Raman spectroscopy is in the important position of being able to clarify the structural characteristics of aqueous electrolyte ~olutions.~-~ This is because the OH stretching, H-0-H bending, and low-frequency Raman bands of water are all sensitive to the structural changes involved for the addition of an electrolyte to water. Numerous infrared and Raman data, from which the structural information about hydrogen bonding in the solutions (1) Lilley, T. H. “Water, A Comprehensive Treatise”;Franks, F. Ed.; Plenum Press: New York, 1973; Vol. 3, Chapter 6. (2) Irish, D. E.; Brooker, M. H. In “Advances in Infrared and Raman Spectroscopy”; Clark, R. J., Hester, R. E., Eds.; Heyden: London, 1976; Vol. 2, Chapter 6. (3) Scherer, J. R. In “Advancesin Infrared and Raman Spectroscopy”; Clark, R. J. H., Hester, R. E., Eds.; Heyden: London, 1978; Vol. 5, Chapter 3.
and vibrational characteristicsbetween water and dissolved ionic species are obtained, have been rep~rted.l-~ In recent pape1-9,~ we have shown that Raman spectroscopy of aqueous solutions in the glassy state can provide an important clue for getting valuable information which is difficult to obtain from Raman spectra of the solutions in the liquid state. The important aspect of Raman spectroscopy for a glassy aqueous solution is the ability to observe low-frequency Raman bands which are obscured and/or hidden by the intense Rayleigh scattering wing in a Raman spectrum for the solution in the liquid state. In addition, the spectral changes associated with going from the liquid state to the glassy state can yield information about the overall effect of a temperature (4) Kanno, H.; Hiraishi, J. Chem. Phys. Lett. 1979, 62, 82; 1980, 72, 541.
0022-3654/83/2087-3664$01.50/00 1983 American Chemical Society
The Journal of Physical Chemistry, Voi. 87, No. 19, 1983 3665
1
Aqueous LiX and CaX, Solutions
&!I
--
Transparent part
Raman
7-
Laser l i g h t
Figure 1. Lower part of the Dewar vessel used for measuring Raman spectra of glassy samples.
change on an aqueous solution and about the geometrical configuration of water molecules around dissolved ions at low temperatures. As an extension and addition of the previous Raman studies of aqueous LiX and CaXz solutions (X = C1, Br, and 1): we here report a detailed study of these solutions with special emphasis on the spectral changes of the OH stretching Raman bands and the low-frequency Raman bands (10-300 cm-l). We also discuss the roles of the charge-transfer states of the OH..-X- bonds and of the local structures around cations and anions in the Raman spectral changes with an addition of electrolyte to water.
Experimental Section Aqueous solutions were prepared by weight by dissolving commercially available (Wako Chemical Co.; guaranteed grade) LiCl, LiBr.H20, LiI.H20, CaCl2-2H20,CaBr2BH20, and Ca12.6H20reagents in distilled water. Vitrification of a sample solution was achieved by immersing an aliquot of the solution in a 4-5 mm inner diameter Pyrex glass cell in liquid nitrogen. The overall cooling rate was about 4 X lo2 K/min. Glass formation was visually checked because transparency of the quenched sample is a good indicator of glass formation, When obtaining a Raman spectrum of a glassy sample, we used a transparent quartz glass vessel designed specifically to maintain the glassy sample solution at liquidnitrogen temperature during the Raman measurements. A schematic diagram is shown in Figure 1. Raman spectra were obtained with a Jasco R-800 spectrometer using 500-600 mW of the 514.5-nm line of a CR-8 argon ion laser as an exciting source. The spectral resolution was 2 cm-' for both the liquid state at room temperature and the glassy state at liquid-nitrogen temperature. The most difficult problem was the elimination of cracks formed in the glassy sample, which can often be a major cause for plasma lines and noise in a Raman spectrum. A commonly employed method for eliminating or reducing the cracks was remelting and revitrifying the sample solution because our experience showed that a crack-free glassy sample may be obtained by chance from several
vitrification runs. Slow or fast cooling of a sample solution was also tried when the above method proved unsatisfactory. For intersample intensity comparison, we adopted a simple internal-standard method in which the v1 band of perchlorate ions added to the sample solution served as an intensity standard. In the previous study,4 we used nitrate ions as the internal standard for the intersample intensity comparison of the low-frequency Raman bands. However, subsequent studies have shown that nitrate ion is unsuitable for the internal standard. A Raman spectrum of a glassy aqueous nitrate solution has a very intense contribution in the frequency region of 10-200 cm-' due to the librational motions of nitrate ions, and thus the nitrate ions added to the halide solutions must have changed to some extent the intensities of the restricted translational Raman bands (10-300 cm-I) of the halide solutions. Therefore, we used instead perchlorate ions as the internal standard because the low-frequency Raman bands are not perturbed by the perchlorate ions to such a large extent as by nitrate ions. In the solutions we added a small amount of perchlorate ion as the internal intensity standard for intersample intensity comparison; the molar ratio H20/C10, was always set to 100 while the ratio of H20/LiX (or CaX2)was set to the respective value ( R = 5,11,12,16, etc.). Here R denotes the molar ratio of water to salt in the solution. The addition of perchlorate ion causes small changes in all Raman bands. However, our rough estimation shows that the contribution to the water Raman spectrum intensities is much less than the molar ratio C104-/X-.
Results and Discussion Effects of Halide Ions on the Raman Band Intensities. The whole range of Raman spectra for the glassy CaX2 solutions (X = C1, Br, and I) are shown in Figure 2. The inset shows a Raman spectrum of the Ca(C104)2solution ( R = 16) to serve as a guide in extracting the intramolecular Raman spectrum for the C104-ions from the Raman spectra for the CaX2 solutions. Similar intensity enhancements of the OH stretching ( Y O H ) and H-O-H bending (JOH) Raman bands in aqueous alkali halide solutions were previously reported by Hornig and his cow o r k e r ~ . More ~ recently, Abe and Ito6 measured the intensity changes quantitatively for aqueous and alcoholic alkali halide solutions in the liquid state at room temperature and revealed the importance of the chargetransfer states in understanding these intensity changes. As cations in these solutions are common, the observed intensity changes are entirely due to anionic effects. Additional evidence for the conclusion comes from the fact that similar intensity changes are also observed in aqueous LiX and HX solutions in both liquid and glassy states though the data are not shown in this paper.7 It is to be noted here that cationic effects, at least among alkali ions, on Raman band (voHand JOHbands) intensities are reported to be very small as compared with anionic effect^.^ When we compare the band intensities by their peak heights, the band is almost twice enlarged in going from the chloride to the iodide solution, the 80H band is about sevenfold enhanced, and the -4O-cm-' band for the iodide solution is about 4.6 times as high as that of the chloride solutions. The observed enhancements for the vOH and &OH bands are essentially of the same magnitudes ~
( 5 ) (a) Busing, W. R.; Hornig, D. F. J. Phys. Chem. 1961, 65, 284. (b) Schultz, J. W.; Hornig, D. F. I b i d . 1961,65, 2131. (6) Abe, N.; Ito, M. J. Raman Spectrosc. 1978, 7, 161. (7) Kanno, H.; Hiraishi, J., unpublished data.
3666
Kanno and Hiraishi
The Journal of Physical Chemistry, Vol. 87,No. 19, 1983
l------l
I I
1100
I
ion
900 600 Raman s h i f t /cm-'
CaI,
30@
soiution
U L i -
CaCl
I
C104
I
i
soiut i o n I
I
R A M A N SHIFT/CM-I Figure 2. Raman spectra of aqueous CaX, solutions (X = CI, Br, and I) in the glassy state at Iiquid-nltrogen temperature. The inset shows a Raman spectrum of perchlorate ions. It is to be noted that all Raman spectra for the CaX, solutions are superimposed by the Raman bands of the perchlorate ions used as the internal intensity standard. The intensities of all Raman spectra were standardized by the u , band (at -935 cm-') of perchlorate ions.
as those reported by Hornig et al.5 and Abe and Ito6 in their Raman intensity measurements on aqueous alkali halide solutions at ordinary temperatures. When seeking to understand the structural and vibrational characteristics of water and aqueous solutions, we immediately consider that the characteristics of hydrogen bonds must play a central role in determining the structural and thermodynamic properties. Before going into a detailed discussion of Raman spectral features of the aqueous LiX and CaX2 solutions, we must point out here that there are two important factors for characterizing spectral features of these solutions though they are intimately correlated and are simply inseparable. The first one is the charge-transfer states of the hydrogen bonds between halide ions and the water molecules coordinated to the halide ions. The second one is the structural changes which involve the breakdown of the basic tetrahedral water structure and the reformation of new local structures around the dissolved ions (cations and anions). In the halide solutions, the inner-sphere hydration around each halide ion consists of about six water molecules,* as depicted in Figure 3, so that the halide ion coordinated by several water molecules can be regarded as a kind of aquo-complex ion and the original hydrogen bonds (OH..-O) of pure water are partly replaced by the hydrogen bonds with halide ions (OH. .X-). Long ago it was reported9J0that halide ions (Cl-, Br-, and I-) dissolved
-
(8) Narten, A. H.; Vaslow, F.; Levy, H. A. J. Chem. Phys. 1973, 58,
5017.
(9)Scheibe, G. 2.Phys. Chem., Abt. B 1929, 5, 355.
Figure 3. Representative configuration of hydration sphere around a chloride ion in aqueous LiCl solution.8 Small closed circles are hydrogen atoms, medium open circles are oxygen atoms, and the large one is a chloride ion.
in water give a strong absorption band in the 200-nm region which arises from the electron-transfer states from a halide ion to the coordinated water molecules through the hydrogen bond (OH...X-) . Abe and Ito6 measured the absolute intensities of the uOH and 60H Raman bands for aqueous alkali halide solutions by changing the concentration of alkali halide and the excitation wavelength (514.5- and 337.1-nm lines) of laser light and concluded that the intensity changes of the VOH and BOH h a n bands are mainly due to the contribution of the charge-transfer states of the [X(HOH),]- complex ions. A remarkable point about all these intensity increases is the difference of the enhancement factors between the uOH and B 0 H bands with the change of halide ion. A similar difference is also observed for the hydrogen bond (OH-. .X-) stretching band at -190 cm-' and its bending band at -40 cm-l. The enhancement factor for the bending Raman bands ( 8 0 and ~ 60~. ,.x-bands) is much larger than that for the stretching Raman bands (uOH and vOH ...Xbands). Molecular orbital calculations made by Webster'l on the molecular geometries of cations and anions in water suggest that in the charge-transfer excited states of the water molecule not only the OH distance but also the HOH angle change considerably from that of the ground-state water molecule. These geometrical and electronic configurational changes in water molecules hydrated to halide ions are undoubtedly the major cause for the difference of the enhancement factors between "OH and 60H (and uOH.. .x- and 60H.. .x-) Raman bands with change of halide ions. However, not enough is known about the molecular charge-transfer states of the hydrogen OH-.-X- bonds to undertake a more quantitative approach. The order of the Raman band intensities in other frequency regions with halide ions (Cl- < Br- < I-) must also be interpreted mainly by the contribution of the chargetransfer states of the OH. .X- bonds. Abe and Ito6showed that the intensity enhancement difference is greatly ex-
-
(10)(a) Fromherz, H.; Menschik, W. Z . Phys. Chem., Abt. E 1929,3, 1; 1930, 7,439. (b) Diamond H.; Fromherz, H. Ibid. 1930, 9,289. (11) Webster, B. J.Phys. Chem. 1975,79,2809. (12)Albrecht, A. C. J . Chem. Phys. 1971,55, 4438.
The Journal of Physical Chemistty, Vol. 87, No. 19, 1983 3667
Aqueous LIX and CaX, Solutions
4000
3000
2000 R A M A N
1600
1200
S H I F T
800
400
0
/cM-'
Flgure 4. Intensity changes of Raman spectrum of aqueous LiCl solution with LiCl concentration. The four black peaks indicated by arrows are due to the intramolecular vibrations of perchlorate ions added as the internal intensity standard.
aggerated by changing the laser exciting line from 514.5 to 337.1 nm. According to the Raman intensity theory,13 the intensity of Raman scattering increases as an exciting line approaches an absorption band of a sample. Knowing that the charge-transfer absorption bands (A- of Cl-, Br-, and I- ions are at -1775 nm for Cl-, -185 and -200 nm for Br-, and -195 and -230 nm for I-,4,l0 we can conclude that the observations by Abe and Itoe are conclusive evidence for the major role of the preresonance Raman effects on the intensity enhancement of Raman bands in aqueous halide solutions with the change of halide from C1- to Brto I-. The other important factor for the intensity changes, which was first pointed out by Hornig et a1.: is the breakdown of the tetrahedral basic water structure and the rearrangements of local configurations around dissolved ions. It is now evident from various Raman and infrared spectroscopic studies1%18that the breakdown and rearrangements of basic water structure by addition of a halide electrolyte are associated with the breaking of intrinsic hydrogen bonds (OH.. .H), collapsing of tetrahedral basic water units, formation of hydration spheres around ions, and resultant hydrogen bonds (OH...X-) . In pure water, Walrafenlg suggested a C2" structure about each water molecule arising from a tetrahedral disposition of oxygen atoms about a central oxygen atom. There have been who claim that there are two types of water molecules depending on the strengths of two hydrogen bonds. In one type, a water molecule has two strong hydrogen bonds, while in the other type a water molecule has one strong and one weak hydrogen bond. The local symmetry of a water molecule in aqueous solu(13) Walrafen, G. E.In "Water,A Comprehensive Treatise"; Franks, Plenum Press: New York, 1972; Vol. 1, Chapter 5. Verrall, R. E. In 'Water, A Comprehensive Treatise"; vol. 3, ,F. Ed.;Plenum Press: New York, 1973; Vol. 3, Chapter 5. Walrafen, G. E. J . Chem. Phys. 1971,55, 768. James, D.W.; Frost, R. L. J . Phys. Chem. 1974, 78, 1754. Subramanian, S.;Fisher, H. F. J. Phys. Chem. 1972, 76, 84. Rodger, G. E.;Plane, R. A. J. Chem. Phys. 1975,63, 818. Walrafen, G. E.J. Chem. Phys. 1964,40, 3249. Murphy, W.F.; Bernstein, H. J. J. Phys. Chem. 1972, 76, 1147. Scherer, J. R.; Go, M. K.; Kint, S. J. Phys. Chem. 1974, 78, 1304. Gorbunov, B.Z.;Naberukhin, Yu. I. Zh. Strukt. Khim. 1975,16, Brink, G.; Falk, M. Can. J. Chem. 1970,48, 3019.
3800
3500
3200
3eoo
R A M A N
3700 S H I F T
3500
3200
2900
'cr-'
Flgure 5. Changes of the OH stretching Raman spectrum from the liquid state at room temperature to the glassy state at liquid-nitrogen temperature. Solid curves are Raman spectra for the glassy state, and dotted ones are those for the liquid state.
tion may also differ from C2"symmetry and is determined primarily by the relative strengths of the two hydrogen bonds. From the geometrical configurations of hydrated spheres of halides and cations, an increase of salt concentration will give rise to an increase in the fraction of water molecules with unequal hydrogen bonds in agreement with the enhancement of the high-wavenumber region in the vOH bands, as discussed later. On the other hand, an increase of strong hydrogen bonds with strong intermolecular coupling intrinsic to water structure causes an intensity increase in the low-wavenumber region (3100 cm-l) in the vOH spectrum. Concentration Dependence of the vOH Raman Spectra in Liquid and Glassy States. Before discussing the vOH Raman spectral features as a function of salt concentration, two Raman spectra (10-4000 cm-l) for the glassy LiCl solutions of different R values ( R = 5 and 11) are shown in Figure 4 to show the general concentration dependence. It is clear from this figure that the boH and boH...x-bands show a larger enhancement than other bands with increase in halide concentration. We believe that the same mechanism as discussed previously is responsible for these enhancement differences. The most interesting feature with which we are mainly concerned here is the changes of the YOH spectral contours with halide concentrataon and with the change of state (namely, from a liquid state to a glassy one). It is a pity that the comparison of low-frequency Raman bands below -300 cm-' in the liquid and glassy states is extremely difficult to make because of the strong Rayleigh scattering wing in a Raman spectrum for the liquid state which prohibits us from observing the lowfrequency Raman bands for the liquid state clearly. As already pointed out, the addition of perchlorate ions as an internal intensity standard causes some changes in the spectral contours in the vOH Raman spectrum and it is very difficult to obtain the absolute intensities of Raman spectra for both liquid and glassy states independently due to the experimental difficulties inherent to the different optical conditions for the liquid and glassy samples. Therefore, we obtained the vOH Raman spectra for both
3668
The Journal of Physical Chemistry, Vol. 87, No. 19, 1983
TABLE I:
Kanno and Hiraishi
Frequencies of t h e Raman Peaks (in c m - ' ) in the Glassy LiX and CaX, Solutions ( X = C1, Br, and I)" solute LiCl LiBr LiI CaC1, CaBr CaI,
VOH
3430 3430 3440 3440 3450 3460
bandb
5 (3440
6 OH
5) i 5 (3450 F 5) 5 (3464 F 5) t 5 (3436 t 5) r 5 (3454 i- 5) i- 5 (3462 F 5) i-
1658 t 1646 i: 1630k 1657 k 1645 r 1624 i
F
3 3 3 3 3 3
bandC (1654 (1644 (1630 (1652 (1643 (1625
U O H . ..X
peakd
200 t 182 i 165 i 200 t 160 F 150i
3) 3) 3) i 3) t 3) i 3) i t t
5 5 5 5 5 5
6 OH.
. . X peake
56 r 42r 33 r 58t 42 t 33r
3 3 3 3 3 3
a Values reported here are improved as compared with the previous ones.4 Values in parentheses are of the liquid state a t The peak freroom temperature. T h e concentration was R = 8 for the LiX solutions and R = 1 6 for the CaX, solutions. guency for the OH stretching Raman spectrum ranging from 3100 to 3700 cm-I. The H-0-H bending Raman band. T h e higher peak for t h e restricted translational latticelike vibrations ranging from 10 t o 300 cm-I. e The lowest frequency peak for the restricted translational latticelike vibrations.
-
liquid and glassy states without an internal intensity standard to see the genuine spectral contours with sacrifice of the intensity changes with salt concentration. The vOH Raman spectra for both liquid and glassy states are presented in Figure 5 as a function of LiCl concentration. A large intensity shift is observed for the spectra of the glassy state with increased salt concentration in contrast to a small change for those of the liquid state at room temperature. In the X-ray diffraction study8 of aqueous LiCl solutions at ordinary temperatures, it has been shown that the basic water structure diminishes with increasing LiCl concentration and is unobserved at R = 10 or less. From the study, the configurations of the Cl(HOH),-, Li(OH2),+, and (H20), ( m N 6, n N 4, and 3c is variable) are suggested to be the most basic structural units in these concentrated LiCl solutions at ordinary temperatures. There have been several report^^,^^,^^ that the frequency of the symmetric OH stretching vibration of the water molecule in liquid water and hydrated crystals is well correlated with the hydrogen-bond strength. A considerable increase in the intensity of the low-frequency region (below 3300 cm-') is observed in the Raman spectrum of amorphous ice26and pure water at low temperatures,2' concomitant with the increase in hydrogen-bond strength and partial recovery in hydrogen bonding with the lowering of the temperature. Thus, the observed enhancement of the low-frequency region in the Raman spectrum for the glassy LiCl solution of R = 10 or 8 clearly demonstrates a partial recovery of the structural configurations intrinsic to the basic water structure which are disrupted by the dissolved ions at room temperature. It is expected that the water-like structure recovered partially in the glassy state also provides the strong intermolecular coupling in the OH stretching vibrations which can be a cause of the intensity increase of the low-frequency region.% This interpretation is partially substantiated by the closeness in the wavenumber of the increased band (peak at -3170 cm-') with the strongest peak at -3100 cm-' from amorphous solid water and polycrystalline ice%where all water molecules are strongly coupled with hydrogen bonds. As the concentration of LiCl increases, the intensity of the low-frequency region becomes weaker in the glassy-state Raman spectrum, indicating the disappearance of water-like structure in the solution. On the other hand, the higher frequency region (peak at about 3450 cm-l) which is mostly due to the OH stretching vibrations of the [X(HOH),]- ions increases in intensity with increasing halide concentration as discussed before. ~~
(24) Novak, A. Struct. Bonding (Berlin) 1974, 18, 177. (25) Joesten, M. D.! Drago, R. S. J.A m . Chem. SOC. 1962, 84, 3817. (26) Li, P. C.; Devlin, J. P. J. Chem. Phys. 1973, 59, 547. (27) Kint, S.; Scherer, J. R. J. Chem. Phys. 1978, 69, 1429. (28) Rice, S. A. Top. Curr. Chem. 1975, 60, 109.
-
-
-
The spectral dip at -3300 cm-' observed in the yoH spectra for the glassy LiCl solutions (R = 8 and 10) is ascribed partly to the result of the Fermi resonance. The importance of Fermi resonance in understanding the OH stretching spectral features has been pointed out by many investigators and the detailed theoretical analysis has been given by Sheref and Sceats et Othen et aL30observed a clear Fermi resonance in the infrared and Raman spectra of polycrystalline tetrachlorocuprate(I1) dihydrate M2CuC14-2H20(M = K, Rb, and NH,) and the dip in the OH stretching spectrum (an Evans hole) significantly becomes deeper on cooling from 300 to 110 K. From these observations, it can be concluded that the effect of the Fermi resonance becomes more important in the vitreous state than in the liquid state in characterizing the vOH spectra. This is clearly evident by the weakness of the dip in the Raman vOH spectra for the liquid state in Figure 5. Frequencies of the Raman Bands. The results of frequency assignments of several Raman bands for the LiX ( R = 8) and CaX2 (R = 16) solutions are given in Table I. As no attempt was made in this study to resolve the broad Raman bands in the OH stretching region into several Gaussian components, we took the peak of the whole envelope as a guide of the strength of the OH stretching vibrations. The peak of the OH stretching Raman spectrum shifts to a higher frequency from the chloride to the bromide to the iodide solution. The OH stretching vibration increases in frequency with decrease in the intermolecular couplings though an unequivocal correlation is not so straightforward for aqueous solutions. As can be seen in Table I, the peak frequency for the OH stretching Raman spectrum is almost unchanged from the liquid state to the glassy state for the calcium halide solutions. On the other hand, it shows a small shift when the lithium salt solutions are vitrified. A peculiar point is that it is almost the same for the glassy LiCl and LiBr solutions (R = 8). We take this coincidence to be rather fortuitous because all other Raman bands behave in the same way as those of the corresponding calcium halide solutions ( R = 16). As shown in Figure 5, the OH stretching Raman spectrum for the lithium halide solutions (R = 8) has two peaks: the higher frequency peak is mainly due to the stretching vibrations of water molecules coordinating to halide ions and the lower one due to the stretching vibrations of water molecuies having configurations similar to the intrinsic water structure. As the OH stretching Raman spectrum is the assembly of several OH stretching vibrational modes arising from the different water configurations, it is very plausible that little difference in the peak frequencies for the glassy LiCl and LiBr solutions is due to the accidental (29) Sceats, M. G.; Stavola, M.; Rice, S. A. J. Chem. Phys. 1979, 71, 983. (30) Othen, D. A,; Knopp, 0.;Falk, M. Can. J. Chem. 1975,53,3837.
The Journal of Physical Chemistry, Vol. 87, No. 19, 1983 3669
Aqueous LIX and CaX, Solutions
combinations of the intensity changes of several Raman envelopes. Except for the OH stretching vibrations, all other Raman peaks seem to move to lower frequency regions with decrease in hydrogen-bond strengths. The librational Raman bands range from about 300 to lo00 cm-I and are apparently insensitive in their spectral contours to the structural changes of the solutions. As already pointed out: the bands move as a whole to lower frequency regions in going from the chloride to the bromide to the iodide solution, and the intensities increase in the same order C1- C Br- < I- as other Raman bands. The Raman bands at -50 and 175 cm-' in pure water at ordinary temperatures are usually ascribed to the hydrogen-bond bending and stretching vibrations, respecti~e1y.l~The -175-cm-l band moves to -213 cm-l in vitreous ice.31 Though the two Raman bands observed below 300 cm-I in the glassy LiX and CaX, solutions must be interpreted as the distributions of the restricted translational latticelike vibration^,^^ we can regard as a simplest picture that the higher frequency band at -180 cm-' (200-150 cm-I depending on anions) is assigned mainly to the stretching vibrations of the OH...O and OH. .X- bonds and the lower frequency one at -40 cm-' (58-30 cm-') to the bending vibrations of the hydrogen bonds. It must be noted here that the frequencies of all Raman peaks are dependent on salt concentration in aqueous solution. This is only natural because the structure and the average hydrogen-bond strength of an aqueous solution change with salt concentration. As there are several types of hydrogen bonds even in pure water, each having a certain distribution in its strength, it is expected that each Raman peak in these solutions actually consists of several components. The two peaks of the restricted translational Raman bands shift to lower frequencies with increasing salt concentration though the shift is small in the glassforming concentration range (R = 4-10 for LiX solution and R = 11-18 for CaX2 solution). Depolarization of the Restricted Translational Raman Bands. The depolarization ratio of a Raman band can give qualitative information about the vibrational nature of the Raman-active vibration. In this study, we measured the depolarization ratios of the restricted translational Raman bands of the glassy solutions and the results are shown in Figure 6. The depolarization ratio for the hydrogen-bond stretching vibration decreases in going from the chloride to the bromide to the iodide solutions. If an oscillator is completely uncoupled, the depolarization ratio for a totally symmetric stretching vibration of the point group Td or 0, should be zero. Therefore, the observed results suggest that the coupling between a halide complex ion [X(HOH),]- (n N 6) and other components (adjacent cations, water molecules, or other halide complex ions) becomes weak in the order C1- > Br- > I-. In this connection, it is to be noted that the property of a halide ion being "a structure breaker" is in the order C1- < Br- C I-.33 Moskovits and MichaelianMmeasured the depolarization ratio for the 170-cm-' stretching and -50-cm-' bending bands of pure water at ordinary temperatures. They gave a value of 0.76 f 0.01 for the depolarization ratio of the
I
h
I
' i
I
I
-
-
-
(31)Hardin, A.H.; Harvey, K. B. Spectrochim. Acta, Part A 1973,29, 1139. (32)(a) Whalley, E.;Berti, J. E. J. Chem. Phys. 1967,46, 1264. (b) Berti, J. E.; Whalley, E. Zbid. 1967,46,1271. (33)Friedman, H. L.;Krishnan, C. V. In 'Water, A Comprehensive Treatise"; Franks, F. Ed.;Plenum Press: New York, 1973;Vol. 3,Chapter 1. (34)Moskovits, M.;Michaelian, K. H. J.Chem. Phys. 1978,69,2306.
R = 16 I
I
r"i
p i . 'I
I
400
I
200
300
1
100 R A M A N
I
300
200
SHIFT
/cM-[
100
I
0
Figure 6. Polarized (11) and depolarized (I) Raman spectra of the low-frequency Raman bands below 400 cm-' for the LiX (R = 8) and CaX, (R = 16) solutions in the glassy state at liquid-nitrogen temperature.
-
170-cm-' band, which indicates the depolarized character of the band. On the other hand, our results gave the values of 0.52 f 0.05 and 0.59 f 0.05 for the - 2 0 0 - ~ m - bands ~ of the glassy LiCl (R = 8) and CaCl, (R = 16) solutions, respectively. Moreover, the depolarization ratios of the 165-cm-l (LiI solution) and 150-cm-' (CaI, solution) bands are nearly zero. These observations, therefore, indicate that the local geometry of a water molecule in these halide solutions is quite different from that in pure water. As aqueous solutions are complicated systems, it is difficult to draw a definite conclusion about the structure only from Raman observations. Weak -280-cm-' Band in the Glassy LiCl and LiBr Solutions. Comparison of the Raman spectra for the glassy Li and CaX, solutions indicates that a weak shoulder near -280 cm-' is observable only for the LiCl and LiBr solutions. The possibility of an observable Li+-OH2stretching Raman band in an aqueous lithium salt solution has long been discussed, and several Raman r e s ~ l t shave ~ ~ ,been ~~ reported, claiming that the LP-OH, stretching Raman band was indeed detected in the low-frequency region: Nash et al.= reported that the totally symmetric vibration of the [Li(OH2),]+ ion was observed at -440 cm-' and Michaelian and M o s k ~ v i t asserted s~~ that the band appeared at 190 cm-' after subtracting the strong Rayleigh
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(35)Nash, C. P.;Donnelly, T. C.; Rock, P. A. J.Solution Chen. 1977, 6,663. (36)Michaelian, K. H.; Moskovits, M. Nature (London) 1978,273,135.
J. Phys. Chem. 1983, 87,3670-3674
3670
scattering wing from the Raman spectrum of a concentrated LiCl solution at room temperature. It is evident that our Raman spectra for the glassy LiX solution support neither the results of Nash et al.35nor the ones of Michaelian and MoskovitsN From the fact that the shoulder at -280 cm-' is not observed in the Raman spectra for the HX3' and CaX, (X = C1, Br, and I) solutions, there can be a possibility that it is associated with some particular configurations of water molecules around Li+ ions. However, the possibility that the shoulder at -280 cm-' may be due to a v1 mode of the [Li(OH,),]+ ions is refuted by the following observations: (1)the shoulder is not clearly a polarized band although the depolarization ratio of the u1 band of the Td symmetry must be zero, and (2) no ob(37)Kanno, H.; Hiraishi, J., unpublished data: Raman spectral features of aqueous HX solutions in the low-frequencyregion (10-300cm-') are similar to those of aqueous CaX2 solutions.
servable Raman band at -280 cm-' is seen in a Raman spectrum for either a glassy aqueous LiI solution ( R = 8) or a LiN03 solution ( R = 8).37 Furthermore, Moskovits and Michaelian have recently reported the presence of a weak, slightly polarized Raman band at -290 cm-' as a component of the restricted-translational Raman spectrum for pure water at room temperaturea3* Thus, it is concluded that the shoulder at -280 cm-' is irrelevant to the ul band of tetrahedral Li(OH2)4+ions. It seems likely that some specific water structure responsible for the -290cm-' band in pure water is retained specifically in the glassy LiCl and LiBr solutions.
Acknowledgment. This work was partially assisted by a Grant-in-Aid for Scientific Research from the Ministry of Education, Science, and Culture of Japan. Registry No. LiC1,7447-41-8;LiBr, 7550-35-8;LiI, 10377-51-2; CaCl,, 10043-52-4; CaBr,, 7789-41-5;CaI,, 10102-68-8.
Perturbation of the Hydrogen-Bond Equilibrium between Cholesterol and Esters or Amides by Chlorofluorocarbons P. Mercler, C. Sandorfy,' D6partement de Chimie, Universltti de Montr&i, Montr&l, Qutibec, H3C 3V1 Canada
and D. Vocelle fipartement de Chimie, Universitti du Outibec 5 Montrtial, C.P. 8888, Montrtial, Qutibec, H3C 3P8 Canada (Received November 1, 1982; In Fhal Form: May 2, 1983)
Infrared spectroscopic studies on hydrogen bonding involving cholesterol show the strong tendency of the latter to self-association. Weaker OH- - -O=C< type H bonds are formed between cholesterol and the carbonyl groups of esters and amides. These are of the same type as the H bonds that cholesterol could form with the carbonyl bonds of phospholipids and sphingolipids. The H-bond equilibria can be perturbed and partly dissociated by chlorofluorocarbon type anesthetics containing an acidic hydrogen. It is believed that this interference with the molecular associationsinvolving cholesterol could be related to fluidization of cell membranes by anesthetics.
Introduction In 1974 Brockerhoff' put forward the idea that cholesterol is hydrogen bonded to proton acceptors present in the hydrogen belt in lipid bilayers. The most likely ones are the ester carbonyl groups of phospholipids and the amide carbonyls of sphingolipids. In Brockerhoff s model the hydrogen belts are well-defined planes of lipid-lipid and lipid-protein H bonding, interposed between the polar heads and the hydrophobic chains of the lipids. Subsequently Huang2q3presented another model in which Hbonding either to the carbonyl group or to water is considered. In their extensive review Demel and de KruyfP made the point that both the hydrophobic and polar parts of the lipid play a role in sterol-lipid interactions and that for this interaction the OH group, planar configuration of the sterol and intact side chain are essential. Through these interactions cholesterol can regulate the permeability and fluidity of biological membranes. (1)H. Brockerhoff, Lipids, 9,645 (1974). (2)C.H. Huang, Lipids, 12,348 (1976). (3)C. H. Huang, Nature (London) 259,242 (1976). (4)R. A. Demel and B. de Kruiff, Biochim.Biophys. Acta, 457,109 (1976). 0022-365418312087-3670$01.50/0
Several infrared and Raman spectroscopic investigations have been made on the associations involving cholesterol."* Parker and Bhaskar5 studied the self-association of cholesterol in CC14and found evidence for the formation of 1:l H-bonded complexes in mixed CCl, solutions of cholesterol and triglycerides. In two extensive papers Bush, Levin, and Levingand Bush, Adams, and Levinloexamined the double-bond stretching region of 1,2-dipalmitoylphosphatidylcholine (DPPC) bilayers and found no evidence for carbonyl-cholesterol H bonds. The downward shift of the carbonyl stretching band was not observed. They concluded that no such H bonds were present. If water was added to the system, the two different carbonyl bands of DPPC coalesced into one broad band without ( 5 ) F. S. Parker and K. R. Bhaskar, Biochemistry, 7, 1286 (1968). (6)R. Faiman, K. Larsson, and D. A. Long, J. Raman Spectrosc., 5 ,
3 (1976).
(7)L. May, C.Baumgartner, and E. R. Cuesta,J. Membrane Biol., 14, 63 (1973). (8) M. Kunst, D. van Duijn, and P. Bordewijk, Z . Naturforsch. A , 34, 369 (1979). (9)S . F.Bush, H. Levin, and I. W. Levin, Chem. Phys. Lipids, 27,101 (1980). (10)S . F. Bush, R. G. Adams, and I. W. Levin, Biochemistry, 19,4429 (1980).
0 1983 American Chemical Society
