J . Phys. Chem. 1993,97, 10622-10629
10622
Infrared Spectra, Relative Stability, and ab Initio Calculations of the CDJCI-HCIvan der Waals Complex Observed in Liquefied Argon W. A. Herrebout and B. J. van der Veken' Department of Chemistry, Universitair Centrum Antwerpen, Groenenborgerlaan 171, 2020 Antwerp, Belgium Received: May 12, 1993'
The mid-infrared (4000-400 cm-I) and far-infrared (300-1 0 cm-I) spectra of several methyl-d3 chloride/ hydrogen chloride mixtures, dissolved in liquefied argon (95-1 17 K) are discussed. In all spectra, experimental evidence was found for the existence of a 1:l van der Waals complex. By use of a temperature study, the complexation enthalpy of the CD3CbHCl complex was determined to be -7.8 f 0.4 kJ mol-'. At higher concentrations of HCl a new absorption band is observed which has to be assigned to a 1:2 bonded complex. Ab initio calculations a t the HF/6-31G** level show that the angular geometry of the CD&l-HCl species is similar to the geometry of the HCl dimer, i.e. that the complex has an approximate L-shape, in which the electron acceptor HCl molecule is bonded through its hydrogen atom to the chlorine atom of the electron donor methyl43 chloride molecule.
Introduction For some time, molecular interactions have been an active field of investigation. Many weak molecular complexes have been studied using molecular beam techniques, combined with rotational spectroscopy and/or high resolution vibrational spectroscopy,' infrared matrix-isolation spectroscopy,2 etc. Some more information about the observed molecular complexes can also be obtained by studying the infrared spectra of the corresponding compounds, dissolved in liquefied noble gases. A more detailed description of several results obtained by using liquefied inert gases (Ar, Kr, Xe, 0 2 , ...) can be found in some recent reviews, published by Tokhadze et al.:, and Kimel'fel'd.4 The observation of weak complexes in liquefied noble gas solutions is strongly facilitated by the low temperatures used and by the high inertness of the surrounding noble gas atoms. In contrast to the low temperature matrices and the molecular beam techniques, the liquid cryogenic systems are characterized by a thermodynamical equilibrium. Therefore, they are well suited for the determination of properties such as relative stability and stoichiometry of the observed complexes. A large number of complexes of the hydrogen halides (HF, HCI, HBr, and HI) and the rather analogous hydrogen cyanide molecules on the one hand and a lot of base molecules, varying fromN2 to (CH3)3N, on the other hand have been studied, mostly by using matrix-isolation spectroscopy and gas phase rotational spectroscopy.s!6 In the literature, several infrared spectra of alkyl halide/hydrogen halide mixtures deposited in an argon matrix have been described,S~'-~but only little information is found about the behavior of those complexes in liquefied noble gases.lOJ1 Therefore, a detailed study of such complexes, formed by different alkyl chlorides and hydrogen chloride, was initiated. In this paper, the results obtained for methyl chloride/hydrogen chloride solutions will be described. To avoid an overcrowding of the 3000-cm-1 region of the spectra, in all cases CDsCl was used. For several CDsCl/HCl mixtures, dissolved in liquefied argon, midinfrared (4000-400 cm-I) and far-infrared (350-1 5 cm-I) spectra were recorded and compared with those obtained for the pure solutions of CD3C1 and HCI, respectively. In addition to the spectra mentioned above, a preliminary structural study of the CH&l.HCI complex was performed using ab initio calculations at the HF/6-31G** level. Before starting the discussion of the spectra of the CD3CI.HCI van der Waals complex, it is interesting to discuss some of the Abstract published in Advance ACS Abstracts, September 15, 1993.
0022-3654/93/2097- 10622$04.00/0
results obtained for analogous complexes, discussed in the literature. The mid-infrared spectra (4000-900 cm-1) of CH3F.HCl and CH:,F.HBr, observed in liquefied argon and krypton, were described by Kolomiitsova et a1.I0 and by Barri et al." Next to the absorption band of free, nonassociated hydrogen chloride, a strong, sharp absorption band near 2799 cm-I was observed in the spectra of liquefied argon containing both CH3F and HCI. This band was assigned to the v(H-Cl) stretching fundamental in the proton donor HCl molecule of the CH3F.HCI complex. Since the hydrogen chloride molecule is hydrogen bonded to the fluorine atom of CH:,F, also small perturbations of the v(C-F) were observed: next to the absorption band near 1041 cm-I, assigned to the v(C-F) in nonassociated CHJF molecules, another absorption band was observed near 1015 cm-I, whose relative intensity strongly increases as the temperature of the studied solutions decreases. This band was assigned as the C-F stretching of the CH3F-HCIcomplex. By use of temperature studies, the complexation enthalpy for the CH3F.HCl molecule was obtained as -4.8 f 0.4 kJ mol-I.lI At higher concentrations of hydrogen chloride, Barri et al." observed another weak absorption band near 2763 cm-I, which was assigned to a higher complex of unknown stoichiometry. Analogous associations between methyl chloride and methyl fluoride on the one hand and hydrogen fluoride on the other hand were also observed in the infrared spectra of solutions in liquefied xenon,3 but because of the low solubility of hydrogen fluoride at temperatures below 210 K, no information about the relative stability could be obtained. Experimental Section The sample of methyl-d:, chloride was synthesized by mixing a small amount of CD3OH (Janssen Chimica, 16.635.48) with PC13 at room temperature. The resulting reaction mixture was separated using a low-pressure, low-temperature fractionation column. The hydrogen chloride was made in small amounts by hydrolyzing PC13 with water and was purified afterward by pumping the reaction mixture through a 2-propanol slush (1 80 K) and followed by fractionation on a low-temperature, lowpressure fractionation column. The argon used has a stated purity of 99.9999% and was therefore used without further purification. All infrared spectra were recorded using a Bruker 1 13v Fourier transform spectrometer. For the mid-infrared (4000-400 cm-I) spectra, a Globar source was used in combination with a Ge/KBr beamsplitter and a liquid nitrogen cooled broad band MCT detector. For the far-infrared (300-10 cm-I) spectra several 0 1993 American Chemical Society
Study of the CD3Cl.HCl Complex
Figure 1. Mid-infrared spectra of pure HC1 (top), pure CD,Cl (middle), and a CD&l/HCl mixture (bottom), dissolved in liquefied argon, at 108
K.
Mylar beamsplitters were used in combination with a liquid helium cooled Si bolometer. The corresponding interferograms of the mid- and far-infrared spectra, recorded at a resolution of 0.5 cm-1, were averaged over 200 and 500 scans, respectively, Happ Genzel apodized, and Fourier transformed using a zero filling factor of 4. The liquid cell, suspended in a vacuum jacket, is made of copper, has a path length of 4.0 cm, and is equipped with wedged Si windows. It is cooled by controlled cold nitrogen bursts and can withstand an internal pressure of 15 bar down to 77 K without leaking. The temperature of the cell is controlled by two PtlOO thermoresistors. The complete cell is connected to a pressure manifold, allowing the filling and evacuation of the cell. After the cell is cooled to the desired temperature, a small amount of the compound(s) is condensed into the cell. Next, the pressure manifold and the cell are pressurized with the noble gas, which immediately starts condensing in the cell, allowing the compounds condensed before to dissolve. In order to get an idea about the solubility of the CD3Cl in liquid argon, the spectra of the liquid argon solutions were compared to that of a crystalline solid film, which was obtained by condensing a small amount of the compound onto a liquid nitrogen cooled CsI window followed by annealing until no further changes were observed in the infrared spectrum. Discussion At the outset, the behavior of the individual compounds was investigated in liquefied argon. Therefore, mid-infrared spectra were recorded of solutions in liquefied argon, containing only hydrogen chloride or methyl-d3 chloride, respectively. A. Hydrogen Chloride. The behavior of hydrogen chloride as a dilute species in an environment of inert noble gas atoms has been the subject of intense vibrational spectroscopic research for many years. Recently, a detailed study of hydrogen chloride dissolved in liquefied noble gases (argon, krypton, and xenon) was described by van der Veken et a1.12 In Figure l a the midinfrared spectrum of hydrogen chloride, dissolved in liquefied argon (108 K), is shown. The dissolved HCl molecules give rise
The Journal of Physical Chemistry, Vol. 97, No. 41, 1993 10623
to a broad, asymmetric absorption band between 3000 and 2700 cm-1, which is determined by the rotational motion of the hydrogen chloride molecules. With an increase of the hydrogen chloride concentration, the infrared spectra start to show a series of weak absorption bands on the low frequency side of the monomer band, that have been assigned1*to hydrogen bonded dimer, trimer, and tetramer hydrogen chloride. B. Methyl-& Chloride. A mid-infrared spectrum of methyld3 chloride dissolved in liquefied argon (108 K) is shown in Figure lb. Next to absorption bands due to CD3 stretching and deformation modes, two relatively intense absorption bands can be observed near 697.5 and 691 cm-I. In the corresponding midinfrared gas phase spectra two intense absorption bands can be observed with well-defined Q-branches near 702 and 695 cm-I, which have been assigned to the C - W l and the C-37C stretching vibration.I3 Since the vibrational frequencies observed in the gas phase spectra and those observed in the corresponding spectra in liquefied argon show only small differences, the absorption bands near 697.5 and 691 cm-l have to be assigned to the C - W l and the C-j7Cl stretches in dissolved methyl-d3 chloride. By use of the simple harmonic approximation, an isotopic splitting of approximate 6 cm-I was calculated for the C-Cl stretching fundamental, which is in very good agreement with the experimental value of 6.5 cm-I. Next to the absorption bands mentioned above, some other, less intense absorption bands are observed near 677, 673, and 668 cm1, the intensities of which strongly increase when the temperature of the liquid argon solution decreases. These absorption bands can be due either to clusters of methyl-d3 chloride, (CDjCl),, or to small particles of crystalline CD3CI. Using molecular beam infrared spectroscopy, Levandier et aI.l4 were able to observe oligomeric species (CH3F),, with n = 2, 3, and 4. The vibrational frequencies observed for these species are situated between the frequencies of the corresponding modes in the monomeric species and those in the CHsF-HCl complex, also investigated by these authors. As will become clear from the discussion below, for methyld3 chloride, in the HC1 complex the C-Cl stretches occur at 689 and 683 cm-I, i.e. above the bands in the 677-668-cm-1 region. Moreover, in thespectraofcrystalline CD3C1, obtained by depositing the compound on a window held at 77 K, bands were observed at exactly the same frequencies. Therefore, it seems unlikely that the bands at 677,673, and 668 cm-1 are due to small clusters, but must be assigned to crystalline particles instead. C. Methyl43 Chloride Hydrogen Chloride Mixtures. The solubility of methyLd3 chloride in liquefied argon increases when a small amount of hydrogen chloride is added to the solutions. A mid-infrared spectrum of such a solution containing approximately 0.8 X 10-3 M CD3Cl and 1.1 X M HCl, recorded at 108 K, is shown in Figure IC. As can be observed in this figure, a new, strong absorption band occurs near 2765 cm-I which is not observed in the mid-infrared spectra of the individual components dissolved in liquefied argon. Therefore it has to be assigned to a van der Waals molecule formed between methyl43 chloride and hydrogen chloride. In Figure 2 the C-Cl stretching region of the spectrum, recorded a t several temperatures between 97 and 116 K, are compared. Next to the absorption bands at 697.5 and 691 cm-I, assigned to the nonassociated CH3CI molecules, two new absorption bands can be observed at 689 and 683 cm-1, whose relative intensity strongly increases when the temperature of the solution decreases. The frequency difference between these absorption bands is of the same magnitude as the isotopic splitting of the C-Cl stretch of the non-associated methyld3chloride molecules and therefore the bands are assigned as the C-35Cl and the C - W l stretch fundamental in associated methyld3 chloride. As can be seen in Figure 2, absorption bands of solute molecules in liquefied noble gases spectra are temperaturesensitive. As the temperature of the solution decreases, the full
+
Herrebout and van der Veken
10624 The Journal of Physical Chemistry, Vol. 97, No. 41, 1993 I
I
1
I
100 I
I
710
700
I
,
680
690 ilcm- I
Figure2. u(C41) region for a CD&I/HCI solution,dissolved in liquefied argon, between 97.6 K (top) and 115.9 K (bottom). width at half height (fwhh) of all absorption bands decreases, while the band maximum is slightly red shifted. The shift of the band maxima can be easily explained by the pronounced temperature dependence of the density of liquid argon.I5 The phenomena described above are similar to those of the CH3F.HCl and the CH3F.HBr complexes, as described by Barri et al.11 Therefore, the structure of the van der Waals complex between CD3Cl and HCl observed here will be comparable to that of the complexes described by Barri et al.; Le. a structure where the hydrogen chloride molecule is hydrogen bonded through its hydrogen atom to the chloride atom of methyl43 chloride. In the mid-infrared spectra studied here, no other experimental evidence was found for the existence of a complex between CD3C1 and HCl, since all other fundamentals assigned to the methyl group do not show any splitting. This of course is not surprising; because of the very weak interactions, only the chemical bond directly involved in the interaction will give rise to pertubed vibrational modes. Also, no new absorption bands were observed that can be correlated with an eventual loss of the C3, symmetry of CD3Cl during the association. At the concentrations of CD3Cl and HCl used to obtain the spectrum in Figure 1c, no bands due to dimeric hydrogen chloride12 can be observed, while it is clear that a considerable amount of CD3CbHCl must be present. The CH3CbHCl complex therefore must be more stable than the hydrogen chloride dimer. When only moderate concentrations (10-3-10-2 M) of CD3C1 and HCl were used, no absorption bands due to a higher complex between both molecules were observed. Because of the low solubility of CD3Cl in liquefied argon, it was impossible to record spectra of concentrated solutions of CD3Cl. However, HCl is rather soluble and therefore relatively concentrated solutions of HCl can be obtained. A part of the mid-infrared spectrum of such a solution, containing approximatly 2.0 X l e 3 M in CD3Cl and 35.0 X 10-3 M in HCI, recorded at 108 K is shown in Figure 3. As can be seen, absorption bands are present at 2770, 2858, 2835, and 2802 cm-I. According to van der Veken et a1.,12 the absorption bands observed above 2800 cm-1 must be assigned to oligomeric hydrogen chloride species. On the low frequency side of the absorption band at 2770 cm-I, assigned to the 1: 1 hydrogen bonded complex between CD3Cl and HCl, another weak absorption band, situated at 27 13 cm-l, is observed;this phenomenon is analogous with the weakabsorption band at 2763 cm-1 observed in the spectra of CH,F/HCl mixtures, dissolved in liquefied argon." Because of the relative weakness of this absorption band in comparison to the other absorption bands in this region, its origin cannot be obtained quantitatively. However, in view of the HCl stretch frequencies of subsequent HCl oligomers, it
1
I
I
I
2900
v.lcm-l
- 1
I
I
I
2700
Figure 3. u(H-CI) region for a 2.0X lo-' M CD&1/35.0 mixture, dissolved in liquefied argon, at 108 K.
X
IO-' M HCI
appears likely that this band has to be assigned to a 1:2 hydrogen bonded complex between CD3Cl and HCl. It is interesting to compare the mid-infrared spectra of the solutions in liquefied argon discussed above with those of the corresponding solid matrices, as described by Barnes.5 In the matrix isolated spectra of CH3Cl/HCl mixtures, a new absorption band was observed near 2745 cm-I which was assigned to the v(H-Cl) stretching fundamental in CHpCl-.HCl. Thus, when passing from the argon solution to the solid argon matrix, the electron acceptor molecule decreases its v(H-Cl) frequency. This is in agreement with the behavior of the HCl oligomers.I2 As for the HCl oligomers, the effect can be correlated to the density change in the environment when passing from liquid to solid. However, in the liquefied argon spectra the absorption band due to the 1:2 species is observed at 27 13 cm-l, which is significantly lower than the frequency of 2730 cm-l, observed in the argon matrix spectra.' For the higher alkyl chlorides, rather analogous phenomena have been observed.16 Relative Stability of the Observed CD$I**.HCI Complex in Liquefied Argon
In order to establish its relative stability, a temperature study of the CD3Cl.HCl complex was performed. Spectra were recorded at temperaturesvarying between 97 and 117 K. At temperatures below 97 K, the greater part of the CD3Cl molecules had crystallized, so only very weak absorption bands due to dissolved CD3Cl and CD3Cl-HCl were observed. Because such spectra do no allow a precise quantitative analysis, they were not used further. The association of CD3Cl and HCl to CD3CbHCl can be described by the equilibrium CD,Cl+ HCl + CD3Cl*HCl and its corresponding stability constant Kq
(1)
In eq 2, CYCD,CI.HC~,C Y C D , ~and ~, CYHC~are the activities for the hydrogen bonded complex, the nonassociated CD3C1, and the nonassociated HCI molecules, respectively. For the concentrations used here, the activities can be safely replaced by the concentrations, which are easily related to relative intensities of bands due to the different species in the experimental spectra. Since it is unlikely that the extinction coefficientswill vary significantly over the concentration range studied here, the formation constant Kq can also be written as
Kq = ICD,CI.HCI 'CD,CI
'HCI
(3)
Study of the CD$l.HCl 1
~
The Journal of Physical Chemistry, Vol. 97, No. 41, 1993 10625
Complex I
I
1
3000
2900
n
I
7
2800
2700
I
2600 */cm-I
720
Figw4. Comparisonoftheexperimentalspectrum (a) with thecalculated fractions of monomer HCI (b) and CD3Cl-HCI (c).
Using the relation AGO = AHo - TASO,this equation can also be written as
RT
(4)
Therefore, plotting h(kD3Cl.HCl/~cD3CIZHCI)versus 1/ T in a Van’t Hoff plot will give rise to a straight line the slope of which is equal to - A H o / R . The asymmetric band due to the free, nonbonded HC1 molecules present in the solutions strongly overlaps with the absorption band near 2768 cm-l, due to the 1:l bonded complex. For the determination of integrated intensities, therefore, the following procedure was adapted. For each temperature at which the CD3Cl/HCl solution was investigated, also the spectrum of a solution only containing HC1 as the solute was recorded. The latter spectrum was rescaled so as to accurately reproduce the HCl monomer contribution above 2800 cm-I in the spectrum of the CD$l/HCl solution. The numerically integrated intensity of the rescaled band was then used as the integrated intensity of the monomer. By subtracting the rescaled spectrum from the spectrum of the CD3Cl/HCI solution, the band due to the complex was isolated, facilitating its integration. The procedure is illustrated in Figure 4. The top spectrum was recorded from a CD3Cl/HCl solution of 108 K; the lower spectrum was obtained at the same temperature from a HCl solution with the same HCl concentration. The middle spectrum is the result of the subtraction. As can be seen in the middle spectrum, a second, weak band is present at 2796 cm-1. The origin of this band will be discussed below. Besides the intensities belonging to the free hydrogen chloride,
700
680 7hn-l
Figure 5. Band profile analysis of the v(C-CI) absorption region of CD3Cl/HCI solutions in liquefied argon, recorded at 99.9 K (A), 106.8 K (B), and 113.5 K (C).
also intensities belonging to the free and associated methyl43 chloride molecules are needed for the construction of the van’t Hoff plot. Both intensities were obtained by least-squares band profile analysis of the experimental spectra in the v(C-Cl) region using Gauss/Lorentz sum profiles. This analysis was hampered by the occurrence in this region of several weak absorption bands due to nondissolved CD3Cl molecules. However, as can be seen in Figure 5 , a good reproduction of the higher frequency region was obtained by using three Gauss/Lorentz profiles, respectively, situated near 678,691, and 689 cm-1. These bands are assigned as the v(C-Cl) stretch fundamental in C D P 2 1 , CD337C1,and CD335Cl.HC1, respectively. By use of the intensities obtained by the methods described above, a van’t Hoff plot was constructed. This plot is shown in Figure 6. From this plot, a complexation enthalpy MO equal to -7.8 f 0.4 kJ mol-1 was calculated. Next to the intensities belonging to the v(H-CI) absorption bands of CD3CbHC1, also the intensity of its v(H-CI) absorption band can be used for determining the corresponding complexation enthalpy AHo. In agreement with the presence of a single type of hydrogen bonded species, the A H 0 value calculated by using these intensities was virtually identical to the one given above. In Table I, the proton affinity as described by B a r n e ~ the ,~ v(H-Cl) stretching frequency in liquefied argon, and the complexation enthalpy of the complex are compared with those for HC1-HCllZand CH3F-HCI.10J1 As described by Arlinghaus et a1.,7 the complexation enthalpy of the complexes can be related to the observed frequency shift of the hydrogen chloride stretching fundamental, Av(H-Cl). Since all complexes in Table I are the result of an electron transfer from the base molecule to the u*
10626 The Journal of Physical Chemistry, Vol. 97, No. 41, 1993
Herrebout and van der Veken
\H
c,
CI............,............. vA\
L/ 1
a i J
H
T
Figure 8. Angular geometry of the CH3CbHCI complex.
I
,
.
.
.
I
.
0.0090
0.0085
.
.
.
l
0.0095
l
l
.
.
l
.
.
.
TABLE 11: Structural Parameters (Bond Lengths in Angstroms, Bond Angles in Degrees, and Dipole Moment in Debye) for CH3CI.HCI Obtained by Using the 6-316** Basis Set
I
0.0100
KT
Figure 6. van't Hoff plot for the CD3CI.HCI complex in liquefied argon. H.
Figure 7. Internal coordinates for the CH3CI.HCI complex.
TABLE I: Properties of Some BaseHCI Hydrogen Bonded Complexes base proton affinitya/ Av(H-CI)/ -AH/ molecule kJ mol-] cm-' kJ mol-] ref HCI CH3F CH3CI a
565 640 669
04 1 070 104
3.8 f 0.4 4.6 f 0.4 7.2 f 0.4
12 11
Taken from ref 5.
orbital of hydrogen chloride, the complexation enthalpy must be correlated with the proton affinity of the base molecules involved. Becauseof thesmaller polarizability of fluorine, the proton affinity of CH3F is much smaller than that of the CH3Cl molecule, so the complexation enthalpy for the CH3F-HCl can be expected to be much smaller than that for the CH3ClmHCl complex. The experimental values for AH', mentioned in Table I, are in line with this prediction. On the basis of the relative intensity of the H-Cl stretching in CD3Cl.HC1, it was predicted above that this complex is substantially more stable than the HCl dimer. This of course is nicely confirmed by the enthalpy differences in Table I. Ab Initio Calculations
Using ab initio calculations the structure for the CH3Cl-HCl complex was predicted. All calculations were carried out using the program BRABO" on a H P 9000/730 workstation. A fully optimized structure for the CH3Cl-.HCl complex was calculated at the H F level, using the 6-31G** split-valence basis set.'* No attempts have been made so far to correct the results for basis set superposition errorlgand/or thermodynamical contributions. In order to be able to interpret the results obtained for the CH3CbHCl complex, analogous calculations were carried out for the individual molecules CH3Cl and HCl, using the same basis set. A. Geometry. By use of the internal coordinates shown in Figure 7, a fully optimized structure for the CH3Cl.HCl complex was calculated by relaxing all internal degrees of freedom without any geometrical constraints. In Table 11, theoptimized geometry parameters for the CH3CbHCl complex are compared with those of the monomers. As can be seen, both the proton donor molecule and the proton acceptor molecule are slightly perturbed when combined into a hydrogen bonded complex. Since the hydrogen bond is formed betwen the hydrogen atom of HCl and the chlorine atom of CH3C1, the bonds involving these atoms slightly increase
r(C 1-C12) r(Cl-H3) r(C 1-H4) r(C 1-H5) r(C12.-H6) r(H6-CI7) L(H3-C 1-C12) L(H4-Cl-CI2) L(H5-C 1-Cl2) L(H3-C 1-H4) L(H3-C 1-H5) L(H4-C 1-H5) L(Cl-C12*.*H6) L(C12.*.H6-C17) L(H3-Cl-C12***H6) L(H4-C I-C12*.*H6) L(H5-C l-C12-.H6) L(C l-C12*..H6-C17) energy (hartree)
dipole moment
CH3CbHCI 1.787 1 .OS3 1 .OS3 1 .OS3 2.784 1.275 108.2 108.3 108.3 110.7 110.7 110.7 99.4 158.6 0.1 120.1 -120.2 3.409 -959.158 160 2.74
HCI
CHaCl 1.782 1.083 1.083 1.083
1.272 108.5 108.5 108.5 110.5 110.5 110.5
-460.0649 18 1.49
-499.090335 2.21
in length upon formation of the complex. Also, the other parameters describing the geometry of the CH3 group hardly are affected by the complexation. The angular geometry of hydrogen bonded dimers can also be described by using the parameters R , CY, p, and T, defined in Figure 8. For the optimized geometry, they were calculated to be 3.99 A, 14.8', 93.S0, and 3.8', respectively. As described by Girardet et al.,zo ab initio calculations at the S C F level slightly underestimate the intermolecular separation since all attractive dispersion forces between both molecules are neglected. However, since the dispersion potential is known to be almost isotropic, the S C F level is in general good enough to determine the angular geometry of a hydrogen bonded species. Since the geometry described above is quite similar to the L-shape found experimentally and theoretically for the hydrogen chloride dimer,z1it can be expected to be close to the real structure, in spite of the relatively low level of approximation of the calculation. The optimized structure of the complex was calculated to have a dihedral angle T I , LH~-C~-C~Z.-H~, close to 0'. Calculations in which T~ is systematically varied, however, show that the barrier to internal rotation is a mere 0.1 kJ mol-I, so that the methyl internal rotation can be regarded as quasi-free; i.e. it is not hindered by the presence of the hydrogen chloride molecule. B. Relative Stability. By use of the calculated energies for CD3C1, HCl, and CDjCl-HCI, the complexation energy can be calculated as
Inthiswayavalueof0.002 907 hartree(7.73 kJmol-1)isobtained. The excellent agreement of this value with the experimental complexation enthalpy, however, has to be regarded as fortuitous: the calculated value was found to be strongly dependent on the basis set and has not been corrected for basis set superposition error,l9 while also the thermodynamic differences between the calculated energy and the 'experimental enthalpy have not been considered. Moreover, it may be noted that the calculated energies are referring to isolated molecules, while the
Study of the CD3ClmHCl Complex
The Journal of Physical Chemistry, Vol. 97, No. 41, 1993 10627
TABLE III: Symmetry Coordinates for CHJCI-HCI aumoximate descriution
svmmetrv coordinate"
CH3 symmetric stretch CH3 asymmetric stretch CH3 asymmetric stretch C-CI stretch CH3 asymmetric deformation CH3 asymmetric deformation CH3 symmetric deformation CH3 rocking CHs rocking CI-H stretch H-CI stretch C-C1-H in plane deformation Cl--H-CI in plane deformation C-CI torsion CI-H torsion a
Not normalized.
thermodynamical measurements described above were all performed in the liquid phase. Although the noble gas solutions have often been considered to be a pseudo gas phase,22.23 recently information has been gathered24 on the influence of the noble gas solvent on some conformational equilibria showing that the thermodynamical equilibria can shift substantially when passing from the vapor phase to the condensed noble gas solutions. C. Vibrational Frequencies. By use of the symmetry coordinates summarized in Table 111, an ab initio harmonic force field for the CHSClaHCl complex was calculated using the techniques described by Pulay et aLzs The force field is given in Table IV. By use of Wilson's FG matrix method,26 from the force field the vibrational frequencies and their potential energy distributions were calculated for CD33SCl.H3sCl. The results are summarized in Table V. Since the C-Cl and the H-Cl bond length increase during the formation of the complex, their calculated stretching frequencies will show a shift to lower frequency. For the C-Cl stretching frequency, the calculated red shift, 11 cm-I, is in good agreement with the experimentally observed shift of 6.5 cm-1. For the v(H-Cl) stretching frequency, the calculated red shift, 23 cm-I, is much smaller than the experimentally observed shift of 104 cm-I. In the calculated frequency shifts of hydrogen halide HX-HY dimers,Z7 similar discrepancies where observed. According to Hannachi et the discrepancies are due to an artifact of the calculations since the neglect of dispersion forces yields an overestimation of the intermolecular separation. It is therefore not surprising that the perturbation felt by the proton donor molecule is underestimated. As described before, the calculated geometry of the methyl group in CD3C1 and CD3Cl.HCl does not shown any pronounced differences. Therefore, no frequency shifts for the corresponding CD3 vibrations are expected, which is confirmed by the experimental results. The small frequency shifts for all CD3 stretching
J 50
I
I
300
250
I
I
I
I
200
150
100
50
VI"'
Figure9. Far-infrared spectra of a pure HCI (a) and CD3CI/HCI mixture (b), dissolved in liquefied argon, recorded at 108 K. The weak shoulder at 220 cm-l in spectrum b is due to an impurity.
fundamentals calculated must therefore be assigned to artifacts of the calculations. Far-Infrared Spectra of CD$3HCl Due to the presence of a weak hydrogen bond in the CD3Cl-HCl complex, low-energy vibrations were calculated at 245.1, 193.5,69.3,33.6, and 14.1 cm-I. In an attempt to observe these vibrations, the spectroscopic study of the solutions was extended into the far-infrared region. Because all the low frequency modes are calculated to be weak, concentrated solutions were used. In agreement with the fact that CD3Cl has no low frequency modes, in the far-infrared spectra of its solutions, no absorption bands due to this compound were observed. In Figure 9, the far-infrared spectrum of a methyl-d3 chloride/ hydrogen chloride mixture, dissolved in liquefied argon (108 K), is compared with that of a pure hydrogen chloride solution. The far-infrared spectra of hydrogen chloride dissolved in liquefied argon have been discussed by Van Aalst et al.28329Due to the hindered rotational motion of the hydrogen chloride molecules, a rotational spectrum can be observed in the far-infrared spectra. Theshouldersobservedat 101,119,142,and 162cm-'in spectrum B of Figure 9 can be assigned to individual rotational transitions,
TABLE IV: Ab Initio (6-31G**) Harmonic Force Constants for CH3CI.HCI SI
s2
s3
s4
ss
6.0091 0.0028 0.0012 -0.1342 0.0004 -0.0002 0.0020 0.0011 -0.0006 0.0005 0.0011 0.0005 0.0001 -0.0004
6.0060 -0.0020 -0.0004 0.1339 0.0004 0.0011 0.0010 0.0010 -0.0010 -0.0018 -0.0010 -0.0001 -0.0003
3.7358 -0.0004 -0.0002 -0.5770 -0.0007 0.0017 0.0853 -0.0095 0.0249 -0.0033 -0.0001 0.0001
0.6283 -0.0004 -0.0003 -0.0318 -0.0018 0.0000 0.0004 -0.0002 0.0010 0.0003 0.0013
s6
s 7
s89
s9
SI0
SI1
SI2
SI3
SI4
SI5
SI 6.0791 S2 273
0.0028 -0.0044 S4 0.1120 S5 -0.0003 s.5 -0,0004 S7 0.1338 0.0012 S9 -0.0020 S i 0 -0.0013 SII 0.0013 Si2 -0.0008 S i 3 -0.0000 Si4 0.0000 Si5 -0.0001
0.6793 -0,0007 0.7420 0.0017 0.0023 0.8329 0.0301 -0.0041 -0.0019 0.8354 -0.0001 0.0062 -0,0004 0.0011 0.0439 0.0005 -0.0022 0.0012 -0.0023 0.0406 5.7828 -0.0001 -0.0086 -0.0033 0.0054 -0.0020 0.0008 0.0591 0.0010 0.0056 0.0015 -0,0026 0.0135 0.0080 -0.1157 0.0207 -0.0002 0.0000 0.0018 0.0010 0.0001 0.0000 -0.0001 0.0002 0.0003 -0.0008 0.0000 -0.0008 -0.0003 0.0000 0.0000 0.0000 0.0000 0.0005 0.0034
Herrebout and van der Veken
10628 The Journal of Physical Chemistry, Vol. 97, No. 41, 1993
TABLE V HCl
Ab Initio (6-31G**) Frequencies (cm-1) and Potential Energy Distribution (PED) for CD33Tl+HCl,CD33TI, and CD335CI*..HCI
frequency 3156.2 2493.1 249 1.9 2312.2 1152.2 1151.9 1132.5
fundamental
PED Sll(99) &(73), S3(26) S3(73), w - 6 ) SI ( 100) S6(83), S s ( w s5(83), ~ ~ ( 1 5 )
CD,)’CI
H-CI stretch CD3 asymmetric stretch CD3 asymmetric stretch CD3 symmetric stretch CD3 asymmetric deformation CD3 asymmetric deformation CD3 symmetric deformation CD, rocking CD3 rocking C-CI stretch C1.-H-CI deformation Cl-H torsion Cl-.H stretch C-CI-H deformation C-Cl torsion
starting from the rotational energy levels with J = 4, 5, 6, and 7, respectively. The artifact at 72 cm-1 is due to the polyethylene windows used in the vacuum jacket. In the far-infrared spectra of a concentrated CH&l/HCl solution, shown in Figure 9b, two new absorption bands are observed, at 283 and 23 1 cm-1. Since these absorption bands do not occur in the spectra of the individual compounds, they are assigned to the vibrations of CH3Cl.HCl calculated at 245 and 193cm-1. The results obtained here are similar to those described by Arlinghaus et ale7:in the far-infrared spectra of argon matrices containing the CH3Cl.HF species, two absorption bands were observed at 436 and 378 cm-I, which were assigned as nondegenerate libration modes of CH3Cl.HF. These latter frequencies are substantially higher than those observed here for CDjClqHCl, which is in agreement with the much stronger hydrogen bond in CHjClmHF. Due to the intense rotational spectrum of hydrogen chloride, the other bands due to the complex cannot be observed in the far-infrared spectra. However, in the mid-infrared region evidence for one of them can be found. It has been observed above that in the subtracted spectrum shown in Figure 4, a second, weak band is present at 2796 cm-I. In the spectra of ethyl-& chloride/ hydrogen chloride mixtures dissolved in liquefied argon, an analogous band was observed in the high frequency wing of the u(H-C1) stretching fundamental belonging to the 1: 1 bonded complex.16 Since the relative intensities in the doublet are not influenced by the temperature or the concentrations used, both bands have to be assigned to the same 1:l bonded complex. A comparable splitting has also been discussed in spectra in liquefied noble gas spectra of several mixtures containing CH3CN and HCL3 In analogy with those results, the weak absorption band at 2796 cm-1 is assigned to the combination band, u(H-CI) + v(C1-H). Since the u(H-Cl) fundamental itself is observed near 2765 cm-I, it follows that the u(C1.-H) stretching frequency must be near 2796 - 2765 = 31 cm-I, which is significantly lower than the calculated value of 69.3 cm-I. D. Interaction Potential. Since the interactions between the molecules observed here are rather weak, one has to take into account that most of thesecomplexes show several large amplitude motions. Therefore, the geometry obtained by using ab initio relaxation techniques, corresponding to a vibrationless structure at 0 K, must be expected to show several differences with the structure that can be determined experimentally. In order to get an idea about these large amplitude motions, an intermolecular potential energy for the CH3Cl.HCl complex was calculated using ab initio techniques. For several values of the intermolecular distance R, ranging form 3.0 to 15.0 A, all other internal parameters were refined completly and the energy of the complex was calculated. The resulting interaction potential is shown in Figure 10.
-10
HC1 3179.3
2486.0 2486.0 2309.9 1152.8 1152.8 1135.0 838.9 838.9 739.8
1 2
4
6
8
10
12
14
16
Figure 10. Intermolecular interaction potential, calculated at the HF/ 6-31G** level.
Due to the rather anharmonic intermolecular potential energy, the intermolecular stretching frequency calculated from the harmonic force field discussed above, will be too high. Therefore, a more accurate stretching frequency was calculated assuming the stretching mode can be described by a one-dimensional Morse oscillator.30 By use of a least-squares refinement procedure the values D, a, and ro of the Morse oscillator were calculated from the ab initio potential to be 585.996 cm-1, 1.244 1/A, and 4.031 A, respectively. From this, the fundamental stretching frequency of the hydrogen bond was calculated to be 43.3 cm-*. This value is in good agreement with the frequency of 3 1 cm-1 derived from the observed u(H.43) + u(C1-H) combination band. Acknowledgment. The N F W O is thanked for financial help toward the spectroscopic equipment used in this study and for a grant to W.A.H. Chris Van Alsenoy is thanked for the ab initio program. Visielab (RUCA) is thanked for the HP 9000/730 computing time. References and Notes (1) Legon, A. C.; Millen, D.J. Chem. Soc. Rev. 1987, 16, 467. (2) Barnes, A. J. In Molecular Interactions; Ratajczak, H . , Orville
Thomas, W. J., Eds.; Wiley: New York, 1980. (3) Tokhadze, K. G.; Tkhorzheskaya, N. A. J . Mol. Srruct. 1992,270, 351.
(4) Kimel’fel’d, Y. M. In Vibrational Spectra and Structure; Durig, J. R., Ed.; Elsevier: Amsterdam, 1992; Vol. 19. ( 5 ) Barnes, A. J. J . Mol. Struct. 1983, 100, 259. (6) Andrews, L. J . Mol. Srruct. 1983, 100, 281. (7) Arlinghaus, R. T.; Andrews, L. J . Phys. Chem. 1987, 91, 1063.
Study of the CD$l.HCl
Complex
(8) Ault, B. S.;Sass, S.E. J . Phys. Chem. 1987, 91, 1063. (9) Barnes, A. J.; Davies, J. B.;Hallam, H. E.;Howells, J. D. R.J. Chem. Soc., Faraday Trans. 2 1973,69, 246. (10) Kolomiitsova, T. D.; Milke, Z.; Tokhadze, K. G.; Schchepkin, D. N. Opt. Spectrosc. (USSR)1979, 46, 391. (1 1) Barri, M. F.; Tokhadze, K. G. Opr. Specrrosc. (USSR) 1981,51,70. (12) Van der Veken, B. J.; De Munck, F. R. J . Chem. Phys. 1992, 97, 3060. (13) Herzberg, G. Infrared and Raman Spectra; D. Van Nostrand Co.: New York, 1966; Vol. 2, Infrared and Raman Spectra of PolyatomicMolecules, p 227. (14) Levandier, D. M.; Mengel, M.; Pursel, R.; McCombie, J.; Scobes. G. Z . Phys. D Ar, Mol, Clusters 1988, 10, 337. (15) Davidson, G.; Davies, C. L. Spectrochim. Acta 1989, 45a, 371. (16) Herrebout, W.A.;vanderVeken,B. J.Tobesubmittedforpublication. (17) Van Alsenoy, C. University of Antwerp, 1992. (18) Francl, M. M.; Pietro, W.J.; Hehre, W. J.; Binkley, J. S.;Gordon, M. S.;Defrees, D. J.; Pople, J. A. J . Chem. Phys. 1982, 77, 3654. (19) Valiron, P.; Vi&, A.; Mayer, I. J . Comput. Chem. 1993, 14, 401,
and references therein.
The Journal of Physical Chemistry, Vol. 97, No. 41, 1993 10629 (20) Girardet, C.; Schriver, A.; Maillard, D. Mol. Phys. 1980, 211,779. (21) Scheiner, S.In Theoretical Treatment oflarge Molecules and Their Interactions; Maksic, Z . B., Ed.; Springer Verlag: Berlin, 1991. (22) Freund, S.M.; Maier, N. B.; Holland, F. R.; Beattie, W. H. J. Chem. Phys. 1978,69, 1961. (23) McLaughin, J. G.; Poliakoff, M.; Turner, J. J. J. Mol. Struct. 1982, 82, 51. (24) Herrebout, W. A,; Van der Veken, B. J. To be submitted for
publica tion. (25) Pulay, P. Mol. Phys. 1969, 17, 197. (26) Wilson. E. B., Jr.; Decius, J. C.; Cross, P. C. Molecular Vibrations; McGraw-Hill: New York, 1955. (27) Hannachi, Y.;Silvi, B. J . Mol. Srrucr. (Theochem) 1989,200,483. (28) Van Aalst, R. M.; Van Der Elsken, J. Chem. Phys. Lett. 1972, 13,
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