4982
J . Phys. Chem. 1986, 90, 4982-4987
Water-Olefin Complexes. A Matrix Isolation Study Anders Engdahl and Bengt Nelander* Division of Thermochemistry, Chemical Center, University of Lund, S-221 00 Lund, Sweden (Received: February 6, 1986; In Final Form: May 1, 1986)
The infrared spectra of the H20, HDO, and D20 complexes with ethylene, propene, cis-2-butene, trans-2-butene, 2methylpropene, 2-methyl-2-butene, 2,3-dimethyl-2-butene, 1,3-butadiene, benzene, and naphthalene have been recorded at temperatures from 11 to 20 K. Interaction energies have been estimated. HDO prefers to form a deuterium bond to the unsaturated molecule; however, small amounts of the hydrogen-bonded forms are present in thermal equilibrium with the deuterium-bonded forms. The energy difference between the hydrogen-bonded and deuterium-bonded forms has been estimated for most of the complexes.
Introduction
The interaction between water and unsaturated hydrocarbons has been the subject of relatively few studies. Karlstrom et ai.' have performed a quantum mechanical calculation of the benzene-water interaction. They found that water prefers to form a hydrogen bond to the benzene 11-electron system; the interaction energy is insensitive to the relative orientations of the interacting molecules near the interaction energy minimum and they could not decide whether a linear or a bifurcated hydrogen bond is preferred. Nilsson2 has recently measured the enthalpies of solution of water in n-alkanes and benzene. He finds a significantly less endothcrmic enthalpy of solution for benzene compared to alkane solutions, supporting the existence of a specific waterbenzene interaction. Several hydrogen halide complexes with unsaturated hydrocarbons have been studied, both with molecular beam-microwave spectro~copy~-'~ and matrix isolation spectros~opy'~-'~ and by ab initio quantum mechanical calculation^.^"-^^ In all cases studied,
(1) Karlstrom, G.: Linse, P.: Wallquist, A,: Jonsson, B. J . Am. Chem. SOC. 1983, 105, 3777. (2) Nilsson, S. 0. J . Chem. Thermodyn., submitted for publication. (3) Read, W. G.: Flygare, W. H.: J . Chem. Phys. 1982, 76, 2238. (4) Legon. A. C.; Aldrich, P. D.: Flygare, W. H. J . Chem. Phys. 1981, 75, 625. (5) Aldrich, P. D.: Kukolich, S. G.: Campbell, E. J. J . Chem. Phys. 1983, 78, 3521. (6) Nelson, D. D.: Fraser, G. T.; Klemperer, W. J . Chem. Phys. 1985. 82, 4483. (7) Shea, J. A.; Bumgarner, R. E.: Henderson, G. J . Chem. Phys. 1984, 80, 4605. (8) Shea, J. A.: Flygare, W. H.; J . Chem. Phys. 1982, 76, 4857. (9) Aldrich, P. D.; Legon. A. C.; Flygare, W. H. J . Chem. Phys. 1981. 75, 2126. (10) Kukolich, S. G.; Read, W. G.: Aldrich, P. D. J . Chem. Phys. 1983, 78, 3552. (1 1) Baiocchi, F. A.: Williams, J. H.: Klemperer, W. J . Phys. Chem. 1983, 87, 2079. (12) Read, W. G.; Campbell, E. J.: Henderson, G . J . Chem. Phys. 1983, 78, 3501. (13) Andrews, L.; Johnson, G. L.; Kelsall, B. J. J . Phys. Chem. 1982, 86, 3374. (14) Andrews, L.: Johnson, G. L. J . Phys. Chem. 1982, 86, 3380. (15) McDonald, S . A.: Johnson, G. L.; Keelan, B. W.; Andrews, L. J . Am. Chem. SOC.1980, 102, 2892. (16) Andrews, L.: Johnson, G. L.: Kelsall, B. J. J . Chem. Phys. 1982, 76, 5767. (17) Andrews, L.; Johnson, G. L.; Kelsall, B. J. J . Am. Chem. SOC.1982, 104. 6180. (18) Andrews, L.: Johnson. G. L.: Davis, S. R. J . Phys. Chem. 1985,89, 1706. (19) Barnes, A. J. J . Mol. Struct. 1983, 100, 259. (20) Sapse, A. M.; Jain, D. C. J . Phys. Chem. 1984, 88. 4970. (21) Hinchliffe, A. Chem. Phys. L e f t . 1982, 85. 531.
0022-3654/86/2090-4982$01.50/0
the hydrogen halide forms a hydrogen bond with a II-orbital on the hydrocarbon. The acetylene and ethylene complexes are T-shaped and the benzene complexes have C, symmetry. This paper extends previous matrix isolation studies of the ethylene-waterZ5 and benzene-water26 complexes to the water complexes of a number of simple olefins and naphthalene. Water forms a hydrogen bond to the II-electron system of the unsaturated hydrocarbon, similar to what has been observed for the hydrogen halide complexes. HDO prefers to form a D bond instead of an H bond, as has been noted p r e v i ~ u s l y . ~ *In - ~ a~ recent study of the hydrogen-bonded complex between water and atomic iodine, we observed the H-bonded HDO-iodine atom complex in thermal equilibrium with the D-bonded complex.32 From the temperature variations of the spectra, we could estimate the energy difference between the two forms to 26 cm-I. The energy difference could be attributed to the zero-point vibration energy difference of one of the bending motions of the hydrogen bond in the two cases. This difference is expected to increase with the strength of the hydrogen bond, and it is therefore understandable that hydrogen-bonded HDO complexes have previously been observed only as metastable forms.28.29~31 The water-hydrocarbon complexes studied here range in strength from the very weak water iodine atom complex up to the water dimer. We have been able to study the equilibrium between the H- and D-bonded forms of the HDO for most of the unsaturated hydrocarbons. The intermolecular vibrations of the water-hydrocarbon complexes differ considerably from those of the water iodine atom complex, so the simple analysis of ref 32 cannot be used unmodified in these cases. However, the energy difference between the H- and D-bonded forms of the HDO complexes appears to increase with the strength of the complex. For the two strongest complexes studied here, the 2-methyl-2butene and 2,3-dimethyl-2-butene complexes, the spectra are independent of temperature in the 11-20 K interval. Experimental Section
The gas mixtures, except with naphthalene, were prepared by standard manometric techniques and sprayed on a CsI window kept at 17 K in a cryostat cooled by an Air Products CS208 refrigeration system. The deposition rate was =9 mmol/h. The deposition system with two separate gas streams has been described.33 (22) Frisch, M. J.; Pople, J. A,: Del Bene, J . E. J . Chem. Phys. 1983, 78, 4063. (23) Hinchliffe, A. Adu. Mol. Relax. Interact. Processes 1982, 24, 245. (24) Pople, J. A.: Frisch, M . J.: Del Bene, J. E. Chem. Phys. Lett. 1982, 91, 185. (25) Engdahl, A,: Nelander, B. Chem. Phys. Lett. 1985, 113, 49. (26) Engdahl, A.; Nelander, 3.J. Phys. Chem. 1985, 89, 2860. (27) Ayers, G. P.: Pullin, A. D. E. Spectrochim. Acta, Part A 1976, 32, 1629, 1641, 1689, 1695. (28) Fredin, L.; Nelander, B.; Ribbegird, J . Chem. Phys. 1977,66,4065, 4073. (29) Nelander, B. J . Chem. Phys. 1980, 72, 77. (30) Nelander, B.: Nord, L. J . Phys. Chem. 1982, 86, 4375. (31) Burneau, A,; Schriver, L.; Manceron, L.; Perchard, J. P. J . Chim. Phys. 1985, 82, 19. (32) Engdahl, A.; Nelander, B. Chem. Phys. 1985, 100. 273.
0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 21, 1986 4983
Matrix Isolation Study of Water-Olefin Complexes
T
0"5
a'L'
3635
3E;Z
2665
3645
?E70
2675
0.45
2680
3.2e
0.:5
3E 5
3620
2632
2655
2664
'L' -2.21 2615
2645
Figure 1. The bound O H ( 0 D ) stretching fundamental of water bound to ethylene. v3(DOH.C2H4) (20 K, upper left), v1(HOD.C2H4) (11 K, upper right), vl(HOH.C2H4)(1 1 K, lower left), and vI(DODC2H4)(1 1 K, lower right). The general shape of the band is best seen in the upper right- and lower left-hand panels. In the lower left-hand panel there is no disturbing water absorption; in the upper right-hand panel the only disturbing absorption, due to D 2 0 , is the small peak to the right of the main peak. The main component of v1(DOD-C2H4)is marked with a C (lower right). The structured absorption to the left of the peak is due to vI(HOD.0D2) (the double peak to the left), v,(HOD.OHD), and vI(HOD.0H2) (the peak with one shoulder on each side). v,(HOD-aq) splits into two components for each isotopomer of water when the temperature is below 17 K. Ar/C2H4 = 168, Ar/total water = 122, 32 mmol of Ar deposited. Note that for the upper left-hand curve we subtracted the 11 K spectrum from the 20 K spectrum in order to straighten the base line. The structure at the left is due to ul(H20-DOH), which shifts slightly with temperature and therefore does not subtract properly. (Absorbance, cm-I).
For the olefins, the concentrations were varied between 1:200 and 1:78. The total concentration of water in the H D O experiments varied between 1:250 and 1:120. A few experiments were carried out with H,O (concentration approximately 1:250). Naphthaleneargon mixtures were prepared by passing pure argon over naphthalene crystals kept at 0 OC. Ethylene (L'Air Liquide 99.95%), benzene (Fisher B255), propene, 2 methylpropene, cis- and trans-2-butene (L'Air Liquid 99%), 1,3-butadiene (L'Air Liquid 99.4%), 2-methyl-2-butene (Fluka, 99.5%), and naphthalene (Baker Pa) were used after degassing. 2,3-Dimethyl-2-butene was purified by gas chromatography on a Perkin-Elmer F21 preparative gas chromatograph with 20% dinonyl phthalate in Chromosorb columns. Water was doubly distilled and degassed. An equilibrium mixture containing approximately equal amounts of H20 and D 2 0 was used to obtain HDO spectra. D2lsO (Miles, 97.5% ISO,99.1% D) was degassed and used without further purification. Argon (L'Air Liquide 99.9995%) was passed through a glass spiral in O,(l). All spectra were run on a Bruker 113v FTIR at 0.5 cm-l resolution. Nomenclature
In complexes of the type studied here, the intramolecular vibrations of the complex-forming molecules are only slightly shifted from their unperturbed positions. Therefore, in order to simplify the notation, the perturbed i-th fundamental of A in a complex with B will be denoted as v,(A.B). C,H, will be used to denote an unspecified hydrocarbon belonging to the set studied here. A water fundamental of H D O which forms an H bond will be denoted a s v,(DOHC,H,) and so on. Results
a. Complex Spectra. When water is codeposited with an olefin in solid argon, there appears a relatively strong band in the symmetric stretching region of the infrared spectrum of water. (33) Fredin, L. Chem. Scr. 1974, 5 , 193.
2656
c
z s i 4/ 3E10
2648
26CQ 0.19
I\
- 0I 0- , 5- - , 3625
0.04
n
T
5576
3584
3502
L.
3500
-~ LEI% ~
2624
2632
Figure 2. The bound O H ( 0 D ) stretching fundamental of water bound to 2-methylpropene. u3(DOH-2-methylpropene) (20 K, upper left), u I (HOD.2 methylpropene) (1 1 K, upper right), vl(HOH.2-methylpropene) (11 K, lower left), and vl(DOD.2-methylpropene) (11 K, lower right) bands. Only in the lower right-hand panel is the band overlapped by other bands. The main components of vl(DOD.2-methylpropene) are marked with a C. The band to the left of vl(DOD.2-methylpropene) is due to u,(DOD.aq). Ar/C4H8 = 81, Ar/total water = 127. 32 mmol of Ar deposited. (Absorbance, cm-I). 0.025
-
i
0.45
nn 3600
3622
3640
3530
3E00
3620
-
I(
26-E
2S23
2669
2640
2680
2660
Figure 3. The bound O H ( 0 D ) stretching fundamental of water bound to cis-2-butene. v3(DOHds-2-butene) (20 K, upper left), vl(HOD.cis2-butene) (1 1 K, upper right), vl(HOH-cis-2-butene) (1 1 K, lower left), and vl(DOD-cis-2-butene) (11 K, lower right). The shape of the band is best seen in the upper right-hand panel where the only disturbing water bands are the two small peaks just below 2660 cm-I, which are due to D 2 0 . In the lower left-hand panel, there is a double peak marked with an A which is due to vl(HOH.0H2). The remaining absorption in this panel is due to vl(HOH-cis-butene). The upper left-hand panel was obtained by subtracting an 11 K spectrum for the same matrix from a 20 K spectrum. In this way we could remove the small components of ul(HOHds-2-butene) which would otherwise appear to the right of the v3(DOH.cis-2-butene)peaks. (Compare the lower left-hand panel.) In the lower right-hand panel, only the two strong components of v I (DODC4H8),marked with C, are seen. The remaining absorption in this panel is due to vl(HOD.aq), vl(HOD.C4H8), and vl(D20). Ar/C4H8 = 78, Ar/total water = 139, 32 mmol of Ar deposited. The lower left-hand panel was taken from an experiment with Ar/C4H, = 162, A r / H 2 0 = 259, 33 mmol of Ar deposited. (Absorbance, cm-I).
The position of the band depends on the olefin; it is shifted toward lower wavenumbers with increasing methyl substitution next to the double bond. With D 2 0 and H D O in the matrix, corresponding bands appear in the symmetric stretching region of D 2 0 and in the OD-stretching region of HDO. The shape of the band does not depend on the water or olefin concentration. In analogy with the previously investigated e t h y l e n e - ~ a t e rand ~ ~ benzenewater26 cases, these bands are assigned to vl(HOHC,H,), v l (DODC,H,), and q(HOD.C,H,) (see Figures 1-4). When HDO is present in the matrix, there appears, in addition to these
4984
The Journal of Physical Chemistry, Vol. 90, No. 21, 1986
I
Engdahl and Nelander
"
A
2
-
Figure 6. Calculated dissociation energies of the water dimer (ref 34-40),
*
:e-2
Z6SO
Figure 4. The bound OH(0D) stretching fundamental of water bound to benzene. u3(DOH.C6H6)(20 K,upper left), uI(HOD.C6H6)( I 1 K, upper right), u,(HOH.C6H6)(11 K,lower left), and Vl(DOD.C,&) (1 1 K, lower right) bands. The only disturbing water absorption in this figure is in the lower right-hand panel, where the absorption below 2640 is due to v,(HOD-water). The vI(HOD.OD2)component is marked with an A. The vl(HOD.OHD) component is barely seen to the right of this peak. All absorption to the right of 2640 is due to v1(DOD-C6H6).Note that the temperature in the lower right-hand panel is 17 K, where the u l (HOD-OD2)has a single peak. Ar/C6H6= 164, Ar/total water = 128, 32 mmol of Ar deposited. The lower right-hand panel is taken from an experiment with Ar/C6H, = 146, Ar/D20 = 150, 31 mmol of Ar de-
posited.
+ 3600
; 3605
3610
3615
3E20
3625
Figure 5. Effect of temperature on u3(DOH-2-methylpropene): upper panel 20 K, middle panel 17 K, lower panel 14 K. (Absorbance,cm-'.)
bands, a band in the OH-stretching region which we did not observe in the previous ethylene-HDO and benzene-HDO experiments. It is quite weak at 11 K but increases in intensity as the temperature increases (Figure 5). At the same time, vl(HODC,H,) and the other bands assigned to D-bonding HDO decrease in intensity. The general appearance of the spectrum of H D O bound to an unsaturated hydrocarbon is very similar to that of HDO bound to atomic where we assigned a similar weak temperature-sensitive band in the OH-stretching region to the OH-stretching fundamental of HDO, which forms an H bond to atomic iodine, v,(DOH.I.). We therefore assign the new HDO band to u,(DOH.C,H,). As can be seen from Figures 1-3, v,(DOH-C,H,) (upper left-hand panel in each figure) has a fine structure very similar to that of vl(HOHC,Hm), v,(HODC,H,), or v,(DODC,H,). (The similarity between u3(DOH.C2H4) and u,(DOD-C,H,) is impossible to see in Figure 1 because of the overlap between v1(DODC2H4)and v1(HOD.0H2), but compare Figure 1 of ref 25, which shows v1(DOD.C2H4)without disturbing HOD.0H2 and HODsOHD adsorption.) In fact, the temperature-sensitive H D O band in the OH region has a fine structure which is similar to vI(HODC,H,) in all cases where we have been
the water formaldehyde complex (ref 41), and the water-benzene complex (ref 1) in kcal/mol vs. vl(HOD.base) in cm-I. The intercept for dissociation energy zero is the vibration rotation origin for v,(HDO) in solid argon from ref 27. able to observe it. v3(DOH*C,&) may be an exception; we are not sure whether the smaller components of v,(HOD-C6H6)are also present for v ~ ( D O H * C ~(Figure H ~ ) 4). For 1,3-Butadiene only the stronger of the two vl(HoD.C4H6) components decreases in intensity with increasing temperature; the weaker component does not change in the 11-20 K interval. At the same time, a broad absorption appears at around 3655 cm-'. For 2-methyl2-butene and 2,3-dimethyl-2-butene, we do not observe any significant temperature effects. We have previously26 found that the fine structure of v I (HOD.C6H6) differs from that of v1(HOH-C6H6)or vl(DOD. C6&) (Figure 4). It is interesting to note that this is an exceptional case; for all other hydrocarbon complexes studied here, the fine structures of v,(HOD.C,H,), vl(HOH.C,H,), and vl(DOD.C,H,) are closely similar. Table I collects the v,(HOH.C,H,), vI(DODC,H,), vl(HODC,H,), and v3(DOH-C,Hm) bands observed in this work. Note that the number of components observed may vary between the bands of the different isotopomers because of overlap with water bands, slight differences in the widths of overlapping components in different regions, or, for v,(DOHC,H,), since some components are too weak to be observable. The remaining fundamentals of complexed water ( H 2 0 , D 2 0 , HDO) are found relatively close to their counterparts in the ethylene25and benzene26complexes and are therefore not given here. Of the hydrocarbon fundamentals, those involving outof-plane bendings of hydrogens next to the double bond are those most affected by the complex formation. They are shifted about 10 cm-l toward higher wavenumbers. For those olefins where the double bond stretching fundamental is active, we observe that it is shifted a few cm-I toward lower wavenumbers upon complex formation. References 25 and 26 give the observed fundamentals of water complexed ethylene and benzene, respectively. b. Dissociation Energies. We have made an attempt to estimate the dissociation energies of the complexes from the observed complex shifts of vI(HODCnHm),assuming that there is a linear correlation between the complex shift and the dissociation energy of the complex. We have based this assumption on a series of HF complexes, for which both calculations and matrix isolation-infrared spectra are available. As long as the dissociation energies are obtained from calculations using large basis sets and involving configuration interaction, the correlation between the measured HF-complex shifts and the calculated dissociation energies is quite good. If one wants to use a similar correlation between measured complex shift and dissociation enzrgy for water complexes, one should use vl(HOD.base) rather than v l (HOH-base) in order to avoid the problems involving the variations in the coupling between the symmetric and antisymmetric O H stretches in complexed water. Except for the water dimer, there are only a very few water complexes for which both complex shifts and calculated dissociation energies are available. Figure 6, which
The Journal of Physical Chemistry, Vol. 90, No. 21, I986
Matrix Isolation Study of Water-Olefin Complexes
TABLE I: v,(HOH.C,H,), vI(DOD.C,H,), vI(HOD-CJIH,),and v,(WHC,,H,)
H0H.B
D0H.B
4985
(Argon Matrix)'
H0D.B
D0D.B
H0H.B
2674.9 2674 2672
2640.8 2639.7 2638.6
2629.4 2628.8 2624.0 2622.7
3613.4 3610.6 3607.8 3605.2 3602.7 3601.0 3597.0 3581.6 3577.9
D0H.B
H0D.B
D0D.B
3600.5 3596.1
2675.5 2672.5 2670.0 2668.4 2665.4 2662 2646.8 2643.7
2620.5 2618.4
2649.7 2640.2 2636.7 2632.0
2622.1 2616.0 2613.1 2610.1
2685.5 2684.1 2681.1 2679.9 2676.1
2646.5 2644.0 2642.7 2641.7
2617.8
2598.3
2678.7 2674.8 2673.9
2644.7 2642.4 2641.3 2640.3
2671.6
2637.8
C2H4
3613.0 361 1.2 3609.5 3606.7
3640.0 3635.8
3595.9 3594.8 3586.5 3584.3
3618.6
2659.8
3606.3 3603.7
2650.8 2649.0
i-C4Hs
C5H10 C3H6
3604.8
3630.0
3596.7
3619.0
2669.5 2667.6 2666.4 2660.4 2659.9
2635.9 2630.6
t-C4H, 3598.9 3596.7 3592.5 3589.0 3588.3 3584.1
2664.6 2661.4 2656.7 2653.7 2652.4 2649.4
3614.9
3584.1 3573.1 3569.0 3563.8
2627.4 2625.8 2625.0 2622.6
C6H6
3620.1 3616.9 3615.9 3612.3 3639.9
C6H12
3570.7 3545.3
C4H6
3618.9
268 1.9
2644.3
-3655 3606.9
2670.7
2636.8
3618.7 3618.0 3614.2 3612.1 3610.9 3610.1 3609.0 3607.7
CIOHS
3634.2
'The data were obtained at 11 K, except for u3(DOH.CnH,), which was measured at 20 K. The q(H0H.B) etc. bands have rather complicated structures (see Figures 1-4). We give the wavenumbers of all components assigned to the band indicated vertically above, with the major component(s) underlined. For comparison, u1(H20) = 3638.0 cm-I, vl(HDO) = 2710.0 cm-I, vl(D20) = 2657.7 cm-l, vl(HOH.0H2) = 3573.6 cm?, vI(HOD-OH2)= 2638.7 cm-I, vI(HOD.OHD) = 2637.6 cm-', vl(HOD.0D2) = 2635.4 cm-I, vI(DOD.0H2) = 2615.9 cm-I, u,(DOD.OHD) = 2615.1 cm-l, and v,(DODO-OD2)= 2613.9 cm-l (ref 27).
TABLE II: Estimated Dissociation Energies (See Text) peak position, dissocn energy, hydrocarbon cm-I kcal /mol 2674.9" 2.4 f 0.5 ethene 2663.8b 3.2 0.6 propene 3.7 0.7 2656.7' trans-2-butene 4.5 f 0.8 2645c cis-2-butene 3.9 f 0.8 2654.4b 2-methylpropene 5.1 f 1.1 2636.7' 2-methyl-2- butene 6.4 f 1.2 2617.8' 2,3-dimethyl-2-butene 2.0 f 0.4 268 1.9' 1,3-butadiene 2.4 f 0.5 2676.1' benzene 2.7 f 0.5 267 1.6' naphthalene
* *
'Main component. bAverage of two almost equal components. Weighted average of the two main components (0.6, 0.4).
we used to estimate the dissociation energies given in Table 11, is based on a series of water dimer calculation^,^^-^^ a waterformaldehyde c a l c ~ l a t i o n , and ~ ' the benzene-water calculation of ref 1. The complex shifts of H D O were taken from ref 27 (34) Pople, J. A. Discuss. Faraday SOC. 1982, 73, 7. (35) Frisch, M. J.; Pople, J. A.; Del Bene, J. E. J . Phys. Chem. 1985.89, 3664. (36) Dierksen, G. H. F.; Kraerner, W. P.; Roos, B. 0. Theor. Chim. Acra 1975, 36, 249. (37) Sokalski, W. A,; Hariharan, P. C.; Kaufman, J. J. J. Phys. Chem. 1983,87, 2803. (38) Clementi, E.; Habitz, P. J . Phys. Chem. 1983, 87, 2815. (39) Newton, M. D.; Kestner, N. R. Chem. Phys. L e r r . 1983, 94, 198. (40) van Lenthe, J. H.; van Dam, T.; van Duijneveldt, F. B.; Kroon-Batenburg, L. M. J. Faraday Symp. Chem. SOC. 1984, 19, 125. (41) Ahlstrom, M.; Jonsson, B.; Karlstrom, G. Mol. Phys. 1979, 38, 1051.
(water dimer), 29 (water-formaldehyde), and 26 (water-benzene). The intercept for zero dissociation energy is 2710 cm-', the value given by Ayers and Pullin2' for the vibration rotation origin of vI(HD0). Since we are unable to decide how accurate the published water dimer dissociation energies are, we have simply drawn three lines, approximately to the lowest, average, and highest of the calculated values. The estimated dissociation energies for the water-olefin complexes given in Table I1 were obtained from the middle line, while the upper and lower lines give the uncertainties. c. H Bonding and D Bonding. The temperature dependency of the observed spectra shows that the lower energy levels of the HDO-hydrocarbon complexes studied here can be divided into two sets, one corresponding to D-bonded and the other to Hbonded HDO. The population ratio between the two sets is given by
n~ = -e-Eo/kT QH n~
QD
where QH and Q D , the partition functions for the H-bonded and D-bonded forms, are given by QH = e-(E,-Eo)lkT, QD = C e-EnIkT iflHI
n4Dl
where Eibelongs to the hydrogen-bonded set of energy levels of the water-d hydrocarbon complex and Eo is the lowest energy level of this set. E, belongs to the deuterium-bonded set of energy levels. The energy levels are given relative to the lowest deuterium-bonded level. Eo is thus the zero-point vibration energy difference between the hydrogen-bonded and deuterium-bonded complexes. Note that the distinction between H and D bonding may not be
4986
The Journal of Physical Chemistry, Vol. 90, No. 21, 1986
TABLE 111: Energy Difference between H- and D-Bonded HDO Complexes (cm-’) from av of from T variation measd ratiosn ethene 19-32 38 propene 19-33 36 trans-2-butene 21-35 >31 cis-2-butene 28-56 39
2-methylpropene benzene naphthalene
31-45 14-2 1 21-25
41 28 28
“See text
meaningful for energy levels 100 cm-I or more above the ground state. The deuterium-bonded form may be changed into the hydrogen-bonded form by a rotation around an axis orthogonal to the water-d plane. The potential barrier separating the Hbonded from the D-bonded minimum along this motion is probably only a few hundred cm-1.32 If we make the assumption that the ratio QH/QDis approximately temperature-independent, we can use the temperaturedependent intensity ratios Z(v3(DOHC,H,))/I(v,(HODC,H,)) to estimate Eo. We first tried to use integrated band intensities, but except for ethylene, the base-line problems in the OHstretching region were too severe and we therefore used ratios of maximum absorbances as estimates of the intensity ratios. Table 111 gives the Eo values obtained from the intensity variations in the 11-20 K inverval. The base-line problems were severe, in particular, at low temperatures and for the stronger complexes, and we therefore give the lowest and highest Eo values consistent with our measurements. We stopped at 20 K in order to minimize the disturbance from water diffusion on the measured values. Since the absorbance per molecule of v,(DOH.C,H,) is approximately twice that of v,(HODC,H,), if the intensity ratio is the same as for free HDO, we can estimate Eo directly from the measured intensity ratio at a given temperature, if we assume Q H = QD.The Eo values obtained in this way appear to vary less than 10% between 11 and 20 K, with the highest values obtained at 20 K. Note, however, that the uncertainties in the low-temperature values are very large (compare Figure 5). Table 111 gives average E,, values for the I1 to 20 K interval.
Discussion Barnes has collected the stretching fundamentals of complexed hydrogen halides obtained from matrix isolation experiments.19 There is a good linear correlation between v(CIHC,H,) and v,(HODC,H,) for the hydrocarbons studied here. The hydrogen chloride complexes of 2-methyl-2-butene and naphthalene do not appear to have been studied. The correlation between vI(HOD-C,H,) and V ( F H . C , H , ) ~ ~ - is ~ *also ~ ~ good, ~ ~ ~ ~with the exception that v( FH.C6H6) seems surprisingly high compared to v,(HOD.C,H,). For those HF complexes where large-scale ab initio calculations have been carried out, there is a linear correlation between the calculated dissociation energies and the complex shifts of H F obtained from matrix isolation experiments. There are some irregularities and the reasons for these are not clear. Nelson et aL6 have recently pointed out that the complex shift of v(FHcyc1opropane) is significantly smaller than one would expect from the induced dipole moment of the complex or from the intermolecular stretching force constant. The difference between observed and expected shift is of the same order of magnitude as the irregularities in the correlation between calculated dissociation energies and observed shifts. If we translate this difference into an uncertainty in the dissociation energy of the cyclopropane-hydrogen fluoride complex via the dissociation energy-complex shift correlation, we get an uncertainty in the dissociation energy of the same order of magnitude as given in Table 11. The example given by Nelson et al. thus serves as a warning: as long as the theoretical connections between the (42) Andrews, I,. J . Phys. Chem. 1984, 88, 2940. (43) Patten, K. 0.;Andrews, L. J . A m . Chem. Sac. 1985, 107, 5594.
Engdahl and Nelander intramolecular complex shifts, intermolecular stretching force constants, etc., and the dissociation energy are unknown, one should not try to make very precise predictions of dissociation energies from measured complex shifts. We believe that the uncertainty limits on the results of Table I1 are reasonably realistic when the above considerations are taken into account. For most of the water complexes studied here, the O H ( 0 D ) fundamental vibration band engaged in the hydrogen bond has a significant fine structure. This fine structure is different for different hydrocarbons, but except for the benzene complex, it is similar for the different isotopmers of a given complex (Figures 1-4). The mechanism behind the fine structure is not clear. If it were due to complexes trapped in different trapping sites, one would expect the intervals between the fine structure components to scale as the band positions between H 2 0 and D 2 0 complexes, as has been found for the water dimer in nitrogen matrices.28 However, the distance between the fine structure components of DzO complexes is around 0.6 of the corresponding H 2 0 spacings, significantly less than the 0.73 expected for complexes trapped in different sites. The ratio between the spacings of the v l (DOD.C,H,) and v,(HOH.C,H,) fine structure components is only slightly larger than the ratio between the largest rotation constants of D 2 0 and H 2 0 , 0.56, perhaps a suggestion that librations of the water part of the complex are involved. In order to investigate the fine structure further, we studied the cis-2butene-D2IRO complex. For this complex, the distance between the two strong components was 2.3 instead of 2.1 cm-I for the D2160complex. The peaks are sharp and easy to measure, so the difference is probably significant. For the DzI80complex, we could observe in addition the smaller fine structure components; the shape of the spectrum was similar to that of HOHecis-2-butene and the spacing of the fine structure components was 0.64-0.65 of the corresponding H 2 0 spacings. The inertia axes of the complexed water molecule are differently oriented relative to the complexed olefin for H- and D-bonded HDO and for the C2,-symmetric water molecules. One may therefore expect the intermolecular motions of a given complex to depend significantly on the isotopomer of water involved. This may be the reason why the fine structure spacings of the H D O complex bands are significantly larger than expected from the H 2 0 and DzO bands (Table I). We are at present unable to offer any model for the intermolecular motions; we can only point to the fact that the spectra of the HzO and D 2 0 complexes are practically independent of temperature in the 11-20 K interval. Unless we believe that the intermolecular dynamics changes significantly when the water molecule is vibrationally excited, we have to assume the presence of one or more excited states within approximately 1-2 cm-l from the ground states for most of the water-hydrocarbon complexes. It is difficult to obtain precise values for the energy difference between the D- and H-bonded forms of the HDO-olefin complexes. If one wants to have measurable concentrations of the H-bonded form over the whole temperature interval, one has to work with rather large water and olefin concentrations. There then appears a broad, somewhat structured absorption in the OH-stretching region, probably due to water bound to olefin aggregates, which makes it difficult to draw a proper base line. The estimates of the energy differences between the D- and H-bonded forms of the complexes are therefore rather uncertain. The estimates obtained at a single temperature from the assumption that QH = QD should be regarded merely as tests of the assumption that the ratio Q H / Q D is approximately temperature-independent. As long as we do not know anything about the relative motions of the water and hydrocarbon molecules, we are unable to say anything about the value of the ratio QH/QDor the intensity ratio between +(DOH-B) and vl(HOD-B). There is no doubt that HDO prefers to form a D bond to olefins, ~~ as it has been found to do to water27i28and f ~ r m a l d e h y d e .There also seems to be a tendency for the energy difference between the H- and D-bonded forms to increase with increasing dissociation energy of the complex. We have not been able to observe the H-bonded forms of the HDO.2-methyl-2-butene and 2,3-di-
4987
J . Phys. Chem. 1986, 90, 4987-4993 methyl-2-butene complexes, for which we estimate dissociation energies equal to or larger than that of the water dimer, an indication that for these complexes the energy difference between the H- and D-bonded forms is at least 60-70 cm-l. For the weakest complexes studied, those of ethylene and HDO and iodine atoms and HDO, the H-bonded form is 20-30 cm-' less stable than the D-bonded form. These results may seem to be in conflict with results from structure d e t e r m i n a t i 0 1 - 1 ,where ~ ~ ~ hydrogen bonds have been found to be shorter than deuterium bonds and therefore generally believed to be stronger than deuterium bonds.47 The conflict is not real, however; the complexes studied here, including the HDO-ammonia complex?0 where only D bonding is observed, are rather weak. For these complexes, the zero-point energy contributions from the intermolecular vibrations determine the relative stabilities of H and D bond^.^*,^^ For strong complexes, the difference in the shifts of the OH- and OD-stretching fundamentals engaged in the bonding will be large and dominate the (44) Gallagher, K. J. In Hydrogen Bonding, Hadzi, D., Ed.; Pergamon: London, 1959. (45) Costain, C.; Srivastava, G. P. J . Chem. Phys. 1964, 4 1 , 1620. (46) Almenningen, A.; Bastiansen, 0.;Motzfeldt, T. Acra Chem. Scand. 1970, 24, 747. (47) Robertson, G. N. Philos. Trans. R. SOC.London 1977, A286, 25.
zero-point vibration energy difference; the hydrogen-bonded form will then 'be more stable than the deuterium-bonded one. Similar considerations should apply to, for instance, equilibria between HX and DX complexes at low temperatures. Here the zero-point energy contribution from the librations of the hydrogen halide should favor the deuterium-bonded complex for weak interactions, while for strong complexes the contribution from the v(XH. B)-v(XD.B) difference will dominate, making the hydrogenbonded complex more stable. When the temperature is raised, the thermal contributions to the equilibrium constants may soon dominate over the zero-point vibration energy contributions. For HDO complexes, the barrier between the H- and D-bonded potential energy minima is probably rather l o ~and~ the~ distinction , ~ ~ between hydrogen bonding and deuterium bonding may therefore cease to be meaningful as the temperature is increased.
Acknowledgment. This work was supported by the Swedish Natural Science Research Council and by Knut and Alice Wallenberg's Foundation. Registry No. H,O, 7732-18-5; HOD, 14940-63-7;D20, 7789-20-0; C6H6, 71-43-2; Ar, 7440-37-1; ethylene, 74-85-1; propene, 115-07-1; cis-2-butene, 590-18-1; trans-2-butene, 624-64-6; 2-methylpropene, 115-1 1-7; 2-methyl-2-butene, 513-35-9; 2,3-dimethyl-2-butene,563-79-1; 1,3-butadiene, 106-99-0;naphthalene, 91-20-3.
Infrared Spectra of Ternary Lithium-Ethylene-Nitrogen Complexes in Solid Argon Laurent Manceron: Michael Hawkins,*and Lester Andrews* Chemistry Department, University of Virginia, Charlottesville, Virginia 22901 (Received: February I O , 1986)
The cocondensation of lithium atoms with ethylene and nitrogen molecules in solid argon produced two kinds of species not seen with lithium and either molecular reagent. The first kind is a series of ternary complexes containing C2H4and N2coordinated to the lithium center and giving rise to N=N stretching and mixed C=C stretching and CH2 deformation modes similar to the binary Li(C2H4)"species. The second species is Li+N2- observed earlier in pure nitrogen, which was not observed in solid argon without a catalytic amount of C2H4present. The presence of C2H4catalyses the fixation of N, in lithium complexes in these experiments.
Introduction The reaction of molecular nitrogen is a long-standing goal of chemistry, owing to the particular stability of the nitrogen molecule and the importance of nitrogen-containing compounds. On the microscopic scale, complexation of metal atoms by nitrogen has been studied in numerous cases with transition metalsId as well as with lithium itself? In the latter case, two new species were reported, the ionic species Li+N2- characterized by an N-N stretching mode at 1800 cm-' (40% reduction in the N-N force constant) and a species containing two nitrogen molecules identified as N2Li,N,. A broad band was also observed at 2300 cm-' and attributed to the N 2 fundamental perturbed by impurity molecules presumably lithium aggregates. However, the production of the two strongly perturbed nitrogen species absolutely required the condensed packing arrangement of the solid nitrogen matrix cage around lithium as these absorptions were not observed in argon matrix studies with small quantities of Li and N2. Likewise, molecular orbital exploration of the Li N 2 energy surface predicted repulsive interaction between Li and N2 in the gas phase because the energy required to populate the Li 2p orbitals is prohibitive; this led the authors to suggest that surface effects or forced approach of Li to N 2 is necessary to overcome
+
f Present address: Laboratorie de Spectrochimie Moleculaire, CNRS VA 508, Bat. F, Universite Pierre et Marie Curie, 5 Place Jussieu, 75230 Cedex Paris, France. *Presentaddress: Sevenoaks School, Sevenoaks, Kent, TN13 lHU, U.K.
0022-3654/86/2090-4987$01.50/0
the energy barrier.7 In this paper we present evidence for an indirect way to obviate the energy barrier necessary for the reaction of molecular nitrogen by preliminary complexation of the metal center to ethylene in solid argon. Experimental Section The cryogenic system, spectrometer, and experimental procedures are identical with that described in the recent Li C2H4 study in this laboratory.* Natural isotopic nitrogen gas (Roberts Co., oxygen free) and two samples of isotopic nitrogen (scrambled, 10% 14-14, 47% 14-15 and 43% 15-15 and mainly I5N2,70% 15-15, 10% 14-15, and 20% 14-14) were used without purification. Isotopic lithium samples (99.99% 7Li and 95.6% 6Li, ORNL) were evaporated from a resistively heated Knudsen cell as described p r e v i o u ~ l y . ~ ~ ' ~
+
(1) Moskovits, M.; Ozin, G. A. J . Chem. Phys. 1973, 58, 1251. (2) Huber, H.; Kundig, E. P.; Moskovits, M.; Ozin. G. A. J . Am. Chem.
SOC.1973, 95, 332.
(3) Ozin, G. A.; Vandervoet, A. Can. J . Chem. 1973, 51, 637. (4) Ozin, G. A.; Vandervoet, A. Can. J . Chem. 1973, 51, 3332. (5) Green, D. W.; Thomas, J.; Gruen, D. M. J . Chem. Phys. 1973, 58,
5453.
(6) Kundig, E. P.; Moskovits, M.; Ozin, G. A. Can. J . Chem. 1973, 51, 27 10. (7) Spiker, Jr., R. C.; Andrews, L.; Trindle, C. J . Am. Chem. SOC.1972, 94, 2401. ( 8 ) Manceron, L.; Andrews, L. J . Phys. Chem., in press. (9) Manceron, L.; Andrews, L. J . Am. Chem. SOC.1985, 107, 563. (10) Andrews, L. J . Chem. Phys. 1968, 48, 972.
0 1986 American Chemical Society
