- Email: [email protected]
The investigation of ion—formamide interactions by ab initio model calculations
Journal of Molecular Structure (Theochem), 124 (1985) 223-230 Elsevier Science Publishers, B.V., Amsterdam - Printed in The Netherlands
THE INVESTIGATION OF ION-FORMAMIDE AB INITIO MODEL CALCULATIONS
INTERACTIONS BY
KAROL MIASKIEWICZ and JOANNA SADLEJ Department
of Chemistry,
University of Warsaw, 02-093
Warsaw (Poland)
(Received 10 December 1984)
ABSTRACT The effect of the cation-molecule and anion-molecule interactions on the frequency shift of the stretching vibrations are discussed in terms of ab initio 4-31 G calculations performed for the model formamide complexes with Li’ and F- ions. The bonding of these systems is discussed qualitatively in the light of the localized orbitals. INTRODUCTION
The vibrational spectroscopy of non-aqueous solutions of electrolytes has proved very successful in investigations of the several basic aspects of the nature of solvent-solute interactions [l] . Among a variety of non-aqueous solvents employed to the spectroscopic investigations of solvent-solute and ion-molecule interactions, formamide seems to be of particular interest [2-61. Understanding of the observed spectroscopic effects is, however, related to the construction of relatively simple models of interacting species. The selected model systems should be small enough for fairly accurate calculations of their properties. Formamide, the simplest model for the peptide linkage, has attracted attention for ab initio calculations on the formamide dimer [7-141, formamide-water complex [8, 9, 15-171 and hydrogen bonding with other systems [ 18-211. The formation and properties of metal ion-formamide complexes have also been studied by Rode et al. [22-261. On the other hand, however, solvation of anions in non-aqueous solutions is gaining interest. The amide-fluoride hydrogen bond has been investigated recently [28]. Possible biochemical implications of this strong hydrogen bond were discussed. The present study is designed to investigate the changes in the electronic structure of formamide molecule and the spectroscopic consequences arising from its interaction with the Li+ cation and F- anion. The simplest pairinteraction model was chosen for this purpose. This model represents a compromise between the complexity of a real solvent-solute system and the simplicity required from the view point of the computations. 0166-1280/85/$03.30
0 1985 Elsevier Science Publishers B.V.
224
Fig. 1. Formamide and its model complexes with ions. COMPUTATIONAL
DETAILS
Ab initio SCF calculations were performed on formamide and its complexes with Li+ cations and F- anions by using the program Monstergauss 81. 4-31 G basis set [29] was applied for H, C, N, 0, F atoms and 5-21 G for the Li atom [ 30 ] . The molecular systems were assumed to be planar according to Fig. l., as is also found experimentally 1311. The calculated parameters are listed in Table 1. The geometry optimization for the LP-HCONH2 complex yields the structure with the cation close to the CO bond axis with Ro.. .I;l+ = 1.69 A, as was found in another paper [ 321. This structure seems to be in agreement TABLE 1 Equilibrium geometry; experimental and calculated, for formamide and its model complex Formamide
4 deg RN
Rco RCH
RNH~ RNH~
RO...w R H... FNC0 NCH CNHt
NCLi+ CNF-
Li+formamide
FormamideF-
4-31G
4-31G
4-31G
4-31G
1.346
1.315
1.321
1.278
1.216
1.245
1.242
1.292
1.081
1.075
1.084
1.076
0.993 0.993 --
0.994
1.115 0.994
w
r311~[401 1.368 1.352 1.212 1.219 1.125 1.098 1.027 1.002 -
125 124.7 112.7 112.7 119.7 120.0 118.7 118.5 -
Li+formamide-F-
124.7
124.3
1.404 127.3
1.386 1.001 1.616 0.983 126.2
113.8
116.0
112.4
115.4
121.9
122.1
120.5
117.4
119.5
120.8
117.6
118.5
-
128.3 -
118.4
125.3 114.2
-
0.997 1.694
-
225
with experimental results of Hertz et al. [33] who measured, for formic acid, the ‘Li relaxation rates in lithium salt solutions. The geometry optimization for HCONH,... F- led to the configuration in which F- is placed on the axis of the NH, bond trans to the CO group with RH . . . F- = 1.40 A. This is in agreement with NMR results [ 341 which indicate the preferential solvation of the Rr- anion by the Pans NH group, as well as with the ab initio calculations [27]. In the case of the complex LitHCONHz .v. F- the geometry was also optimized. The distances of 0. eeLi+ and H.. - F- are a little shorter than in the single complexes. Relative to the parent formamide, the CN bond is shortened and CO bond is lengthened. The R H...F -distance is almost the same as in HF molecule, RHF,exp.= 0.917'A,while the NH, bond is lengthened about 0.4 A. Force constants were calculated by double numerical derivation of the electronic energy of the system. The energy dependence on the bond length was approximated by the polynomial in the Monstergauss 81 program. RESULTS AND DISCUSSION
The results of ab initio 4-31 G energy calculations are presented in Table 2. The hydrogen bond between formamide molecule and fluoride ion belongs to the very strong asymmetric hydrogen bonds, Emsley et al. [28] propose to define the energy of this hydrogen bond with respect to the amido anion and hydrogen fluoride molecule instead of the amide molecule and fluoride ion. However, we compute the energy of the hydrogen bond with F- in a manner similar to that used in the energy calculations for the cation-molecule complex, i.e.j with respect to the formamide molecule and fluoride anion. The interaction energy is as follows: for Li+-formamide complex -62.3 kcal mol-I, for F--formamide system -51.5 kcal mol-’ and for the system Li+-formamide-F-205.8 kcal mol-‘. The calculated force constants are listed in Table 3. It is seen from Table 3 that all the complexes show considerable changes in force constants of the three stretching vibrations of the CO, CH and NH groups. The cation slightly decreases the CO frequency and increases the CH frequency. On the other hand, the anion decreases the CO as well as NH and CH frequencies. TABLE 2 Results of the energy 4-31G
calculations on the formamide complexes with ions
Species
Total energy (a.u.)
Formamide Li+-formamide F--formsmide Li+-formamide-FFluoride ion Lithium cation
-168.681588 -176.015114 -268.011430 -275.490751 -99.247824 -7.233262
226 TABLE 3 Force constants for the stretching vibrations in mdyn a-’ Species
kc0
kcri
kN?x
Formamide Li+-formamide F--formamide
13.8 13.6 11.8
5.80 6.07 5.77
Li+-formamide-F-
12.2
6.05
8.15 8.15 4.99 8.07 1.72 7.81 3.98
(sym NH,) (sym NH,) (sym NH,) (NH,) (NH,) (NH,) (NHt)
The lower CO bond strength in the lithium-cation complex and fluorideanion complex is also reflected in the evaluated equilibrium bond distances Rco. The CO bond on all the complexes are slightly longer than that in the formamide. molecule. Conversely, the CN bonds are slightly shorter. It should be pointed out that Ottersen and Jensen [Xi] and Ottersen et al. [ 181 obtained different results for the formamide-water [ 151 and formamidewater [ 15) and formamideammonia [ 181 systems. In these papers the kc0 values increase upon formation of the hydrogen bond. This result is rather puzzling. For the NH bond we calculated the force constants for the symmetric vibration and for the stretching one of the NH bond in the case of the complexes with anions. Their values are very different. It should be noted that both these quantities, the stretching force constants and the bond lengths, correspond only to the net effect of the rearrangement of electron density upon complex formation. However, it is interesting to look for the more detailed analysis of the electron density redistribution in the perturbed system. It can partly be achieved by means of an analysis of the atomic population. Table 4 presents the net atomic population for the investigated molecular systems. TABLE 4 The net atomic charges for foramide and its complexes Atom
Formamide
Lit-formamide
F--formamide
Li+-formamide-F-
C
0.587 -0.904 -0.612 0.170 0.373 0.386 0.0
0.700 -0.855 -0.831 0.261 0.428 0.428 0.869 0.123
0.565 -0.962 -0.740 0.133 0.521 0.309 0.825 -0.175
0.597 -0.868 -0.934 0.195 0.585 0.320 0.779 -0.676 -0.103
N 0 H Ht HC Li+ FTotal net charges in formamide
227
When comparing the net atomic populations in complexes with those in the parent formamide molecule, a considerable charge redistribution is noticed. An examination of the changes indicates that there is a charge transfer from anion to formamide molecule. In the complex with cations, however, electron density is transferred in the opposite direction. The increase in the polarity of the CN and CO region is observed: the oxygen atom gains electrons in all complexes, while the changes on the nitrogen and carbon atoms are in opposite directions upon the coordination of cation and anion. This picture of electron density changes in the formamide molecule, which results from the adduct formation, is more directly obtainable from the set of localized molecular orbit&. The localized MO are examined according to Boys [41]. In Table 5 we present the coordinates of the centroid of charge with respect to the molecular coordinate system for all the localized molecular orbit&. In the complex with a cation, the centre of charge for two CO LMO is shifted towards the oxygen atom, for two NH LMO towards the nitrogen atom. In the complex with an anion, one LMO which previously described the NH, bond is now identified, using critical level 0.080, as centered on the N atom. The complex OCH-NH2 - F- seems now to be the HOCNH-. . . H’. . . F- system. An interesting effect is observed when two ions are placed close to the parent formamide molecule. Now, one of the LMOs is centered on the N atom, but one of the LMOs which described the CO bond is centered only on the oxygen atom. It induces a change in the hybrydization of oxygen lone pairs, which now have the centre of charge located out of the molecular plane and the CO bond seems to be a single bond rather than a double one: Li’...HO---C--NH~o.H’.*.F-. Finally, the results presented above can be compared with some experimental observations in the Raman spectra. The model calculations refer to the l
l
TABLE 5 Centroid of charge (x, y, z in bohr) for formamide and its complexes LMO
formamide
Li+-formamide
F--formamide
Li+-formamide--F
ls0 ISN ISC NH+ NH; co co CN CN CH
1.89 0.00-1.31 0.00 0.00 0.00 0.00 0.00 0.00 -0.94 0.00 3.13 0.95 0.00 3.09 1.23 0.46-0.85 1.23 -0.46-0.85 0.00 -0.50 1.87 0.00 0.50 1.87 -1.28 0.00 -a.54 2.38 0.00 -0.99 1.75 0.00 -1.88 -
1.94 0.00 -1.32 0.00 0.00 2.48 0.00 0.00 0.00 -0.92 0.00 3.05 0.93 0.00 3.03 1.39 -0.44 -0.94 1.39 0.44 -0.94 0.00 0.51 1.75 0.00 -0.51 1.75 -1.20 0.00 -0.57 2.50 0.00 -1.03 1.85 0.00 -1.96 4.36 0.00 -3.44 -
1.87 0.00 -1.42 0.00 0.00 2.49 0.00 0.00 0.00 -0.81 0.00 3.0W 1.01 0.00 3.02 1.27-0.45~.97 1.27 0.45 -0.97 0.01-0.51 1.75 0.01 0.51 1.75 -1.30 0.00 -0.54 2.38 0.00-1.13 1.71 0.00-2.01 -4.17 0.00 4.75 -3.58 0.00 4.45 -4.47 -0.41 4.57 4.41 0.45 4.63 4.09 0.00 5.27
1.97 O.OO-1.44 0.00 0.00 2.41 0.00 0.00 0.00 -0.69 0.00 2.83a 1.01 0.00 2.93 1.18 0.04 -0.88 2.07-0.53-l.32b 0.02 0.53 1.53 0.02 AI.53 1.53 -1.22 0.00-0.61 2.39 0.37 -1.16 1.82 0.12 -2.07 4.49 0.00-3.17 4.07 0.00 4.25 -3.37 0.00 4.00 4.16 0.46 4.49 4.38 4.06 3.83 4.13 -0.40 4.57
IPO lP0
1sLi ISF 2sF 2pF 2pF 2pF
aLMO is centered on the N atom. bLMO is centered on the 0 atom.
228
isolated molecule and its isolated complex with cations and anions. These data should be compared rather with the gas-phase spectra of formamide. The only experimental data available at that moment are for the liquid solvent and the corresponding solutions of electrolytes. However, we can try to use the information following from the comparison of the gaseous and pure liquid formamide. Amide I band frequency It is proposed that, in the liquid phase, the formamide molecules tend to form, most probably, antiparallel, cyclic dimers [3]. It was found that the CO/amide I band frequency is 95 cm -’ lower in pure liquid formamide than in the gas phase [ 351. This latter observation suggests that the breaking of hydrogen bonds between self-associated formamide molecules in gas phase produces the amide I frequency increase. This is consistent with ab initio SCF calculations for hydrogen bonded cyclic dimer [ 121. When electrolytes are dissolved in formamide the frequencies of the amide I band are higher than in the pure liquid formamide [ 4,6]. However, in the ternary solutions, formamide + acetonitrile + salt, [6] for various salts the results show a decrease in the amide I band frequency. According to the authors [6, 391, these results prove that cations decrease the amide I frequency being linked to the CO group, while the anions increase the frequency changes provoked by cation. This conclusion comes from the assumption that in the binary solutions, formamide + acetonitrile, some of formamide dimers are broken by the formamide-acetonitrile interaction and, next, upon the dissolution of the electrolyte, the preferential solvation of ions by the amide can be found [3]. Theoretical results for the model systems presented in Table 3 show the small decrease of the CO force constant after the coordination of the cation through the CO bond and the rather drastic decrease after the interaction with fluoride anion. It should be kept in mind, however, that the CO stretching contributes only 65% to the amide I mode [36] or even less [38]. The rest is CN stretching vibration, 6 NH, and VN-n, which also changes after the coordination with a cation or anion. In spite of the approximation, assuming that the CO bond of the formamide molecule vibrates independently, there is a qualitative agreement between the computed frequency changes and those found in the spectra [6]. The CH stretching band The CH stretching band is influenced by ions in a different manner. As follows from Table 3, a cation causes an increase in its frequency but an anion decreases it. Nevertheless, there are no other calculations on the CH stretching frequency for the model system. Thus, we can compare our results obtained for the isolated system with the experimental data for solutions of
229
electrolytes in HCOND2. Raman spectra of the binary and ternary solutions of electrolytes indicate an increase of the CH stretching band caused by cations and a decrease caused by anions. Because the Yc_n stretching vibration is very characteristic [36], we can say, that these findings are in agreement with the calculations. The NH stretching band Both ions cause a decrease of the NH stretching band for the isolated system. The experimental results [4] indicate the lowering of the ND stretching frequency of the HCONHD upon dissolved the salts. Because both vibrations are characteristic [ 37, 381, also a lowering VN-HS~~. andv~--H~~~. of the symmetric NH2 stretching frequency is expected, in agreement with the presented calculations. ACKNOWLEDGEMENTS
We are indebted to Professor Z. Kecki for his interest and helpful comments on the manuscript. This work was supported, in part, by the Polish Academy of Sciences in connection with project MR.I.9. REFERENCES 1 D. E. Iresh and M. H. Brooker, in R. J. H. Clarke and R. E. Hester (Eds.), Advances in Infrared and Raman Spectroscopy, Vol. 2, Heyden, London, 19’76. 2 M. H. Baron and C. de Loze, J. Chim. Phys. Phys. Chim. Biol., 69 (1972) 1084. 3 M. H. Baron, H. Jaeschke, R. M. Moravie, C. de Loze and J. Corset, in Metal-Ligand Interactions in Organic Chemistry and Biochemistry, Part 1, D. Reidel, Dordrecht, 1977, p. 171. 4 J. Bukowska, Chem. Phys. Lett., 57 (1978) 624. 5 D. J. Gardiner, R. B. Girling and R. E. Hester, J. Chem. Sot. Faraday Trans. 2, 71 (1975) 709. 6 J. Bukowska and K. Miaskiewicz, J. Mol. Struct., 74 (1981) 1. 7 M. Dreyfus and A. Pullman, Theor. Chim. Acta, 19 (1970) 20. 8 A. Pulhnan, H. Berthold, C. Giessner-Prettre, J. F. Hinton and R. D. Harpool, J. Am. Chem. Sot., 100 (1978) 3991. 9 J. F. Hinton and R. D. Harpool, J. Am. Chem. Sot., 99 (1977) 349. 10 S. Yamabe, K. Kitaura and K. Nishimoto, Theor. Chim. Acta, 47 (1978) 111. 11 T. Ottersen and H. H. Jensen, J. Mol. Struct., 26 (1975) 355. 12 T. Ottersen, J. Mol. Struct., 26 (1975) 365. 13 F. A. Momany, R. F. McGuire, J. F. Yan and H. A. Scheraga, J. Phys. Chem., 74 (1970) 2424. 14 D. Peters and J. Peters, Int. J. Quantum Chem., 19 (1981) 1121. 15 T. Ottersen and H. H. Jensen, J. Mol. Struct., 26 (1975) 375; 28 (1975) 220. 16 J. E. Del Bene, J. Chem. Phys., 62 (1975) 1314,1961. 17 J. E. Del Bene, J. Am. Chem. Sot., 100 (1978) 1387,1395. 18 T. Ottersen, H. H. Jensen, R. Johansen and E. Wisloff-Nilssen, J. Mol. Struct., 30 (19760 379. 19 P. Otto, S. Suhai and J. Ladik, Int. J. Quantum Chem. Biol. Symp., 4 (1977) 451. 20 R. Janoschek, Theor. Chim. Acta, 32 (1973) 49.
230 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41
R. Janoschek, The Jerusalem Symposia, 6 (1974) 284. B. M. Rode and H. W. Preuss, Theor. Chim. Acta, 35 (1974) 369. R. Fusseneger and B. M. Rode, Chem. Phys. Lett., 44 (1976) 95. B. M. Rode and,R. Fusseneger, J. Chem. Sot. Faraday Trans. 2,71 (1975) 1958. D. N. Fuchs and B. M. Rode, Chem. Phys. Lett., 82 (1981) 517. B. M. Rode, in Metal-Ligand Interactions in Organic Chemistry and Biochemistry, Part 1, B. Reidel, Dordrecht, 1977, p. 127. P. Schuster, in L. Kevan and B. Webster (Eds.), Electron-Solvent and Anion-Solvent Interactions, Elsevier, Amsterdam, 1976, p. 259. J. Emdey, D. J. Jones, J. M. Miller, R. E. Overill and R. A. Waddilove, J. Am. Chem. Sot., 103 (1981) 24. R. Ditchfield, W. J. Hehre and J. A. Pople, J. Chem. Phys., 54 (1971) 724. P. C. Hariharan and J. A. Pople, Mol. Phys., 27 (1974) 209. M. Kitano and K. Kuchitsu, Bull. Chem. Sot. Jpn, 47 (1974) 67. P. Shuster, W. Jakubtz, W. Marius, Top. Current Chem., 60 (1975) 1. H. G. Hertz, H. Weingartner and B. M. Rode, Ber. Bunsenges. Phys. Chem., 79 (1975) 1190. R. D. Green, Can. J. Chem., 47 (1969) 2497. S. T. King, J. Phys. Chem., 75 (1971) 405. I. Suzuki, Bull. Chem. Sot. Jpn, 33‘(1969) 1359. Y. Sugawara, Y. Hamada, A. Y. Hirakawa, M. Tsuboi, S. Kato and K. Morokuma, Chem. Phys., 50 (1980) 105. E. D. Schmid, E. Brodbeck, J. Chem. Phys., 78 (1983) 1117. J. Bukowska, J. Mol. Struct., 98 (1983) 1. E. Hirota, R. Sugisaki, C. J. Nielsen and G. 0. Sorensen, J. Mol. Spectrosc., 49 (1974) 251. S. R. Boys, Rev. Mod. Phys., 32 (1960) 306.
THE INVESTIGATION OF ION-FORMAMIDE AB INITIO MODEL CALCULATIONS
INTERACTIONS BY
KAROL MIASKIEWICZ and JOANNA SADLEJ Department
of Chemistry,
University of Warsaw, 02-093
Warsaw (Poland)
(Received 10 December 1984)
ABSTRACT The effect of the cation-molecule and anion-molecule interactions on the frequency shift of the stretching vibrations are discussed in terms of ab initio 4-31 G calculations performed for the model formamide complexes with Li’ and F- ions. The bonding of these systems is discussed qualitatively in the light of the localized orbitals. INTRODUCTION
The vibrational spectroscopy of non-aqueous solutions of electrolytes has proved very successful in investigations of the several basic aspects of the nature of solvent-solute interactions [l] . Among a variety of non-aqueous solvents employed to the spectroscopic investigations of solvent-solute and ion-molecule interactions, formamide seems to be of particular interest [2-61. Understanding of the observed spectroscopic effects is, however, related to the construction of relatively simple models of interacting species. The selected model systems should be small enough for fairly accurate calculations of their properties. Formamide, the simplest model for the peptide linkage, has attracted attention for ab initio calculations on the formamide dimer [7-141, formamide-water complex [8, 9, 15-171 and hydrogen bonding with other systems [ 18-211. The formation and properties of metal ion-formamide complexes have also been studied by Rode et al. [22-261. On the other hand, however, solvation of anions in non-aqueous solutions is gaining interest. The amide-fluoride hydrogen bond has been investigated recently [28]. Possible biochemical implications of this strong hydrogen bond were discussed. The present study is designed to investigate the changes in the electronic structure of formamide molecule and the spectroscopic consequences arising from its interaction with the Li+ cation and F- anion. The simplest pairinteraction model was chosen for this purpose. This model represents a compromise between the complexity of a real solvent-solute system and the simplicity required from the view point of the computations. 0166-1280/85/$03.30
0 1985 Elsevier Science Publishers B.V.
224
Fig. 1. Formamide and its model complexes with ions. COMPUTATIONAL
DETAILS
Ab initio SCF calculations were performed on formamide and its complexes with Li+ cations and F- anions by using the program Monstergauss 81. 4-31 G basis set [29] was applied for H, C, N, 0, F atoms and 5-21 G for the Li atom [ 30 ] . The molecular systems were assumed to be planar according to Fig. l., as is also found experimentally 1311. The calculated parameters are listed in Table 1. The geometry optimization for the LP-HCONH2 complex yields the structure with the cation close to the CO bond axis with Ro.. .I;l+ = 1.69 A, as was found in another paper [ 321. This structure seems to be in agreement TABLE 1 Equilibrium geometry; experimental and calculated, for formamide and its model complex Formamide
4 deg RN
Rco RCH
RNH~ RNH~
RO...w R H... FNC0 NCH CNHt
NCLi+ CNF-
Li+formamide
FormamideF-
4-31G
4-31G
4-31G
4-31G
1.346
1.315
1.321
1.278
1.216
1.245
1.242
1.292
1.081
1.075
1.084
1.076
0.993 0.993 --
0.994
1.115 0.994
w
r311~[401 1.368 1.352 1.212 1.219 1.125 1.098 1.027 1.002 -
125 124.7 112.7 112.7 119.7 120.0 118.7 118.5 -
Li+formamide-F-
124.7
124.3
1.404 127.3
1.386 1.001 1.616 0.983 126.2
113.8
116.0
112.4
115.4
121.9
122.1
120.5
117.4
119.5
120.8
117.6
118.5
-
128.3 -
118.4
125.3 114.2
-
0.997 1.694
-
225
with experimental results of Hertz et al. [33] who measured, for formic acid, the ‘Li relaxation rates in lithium salt solutions. The geometry optimization for HCONH,... F- led to the configuration in which F- is placed on the axis of the NH, bond trans to the CO group with RH . . . F- = 1.40 A. This is in agreement with NMR results [ 341 which indicate the preferential solvation of the Rr- anion by the Pans NH group, as well as with the ab initio calculations [27]. In the case of the complex LitHCONHz .v. F- the geometry was also optimized. The distances of 0. eeLi+ and H.. - F- are a little shorter than in the single complexes. Relative to the parent formamide, the CN bond is shortened and CO bond is lengthened. The R H...F -distance is almost the same as in HF molecule, RHF,exp.= 0.917'A,while the NH, bond is lengthened about 0.4 A. Force constants were calculated by double numerical derivation of the electronic energy of the system. The energy dependence on the bond length was approximated by the polynomial in the Monstergauss 81 program. RESULTS AND DISCUSSION
The results of ab initio 4-31 G energy calculations are presented in Table 2. The hydrogen bond between formamide molecule and fluoride ion belongs to the very strong asymmetric hydrogen bonds, Emsley et al. [28] propose to define the energy of this hydrogen bond with respect to the amido anion and hydrogen fluoride molecule instead of the amide molecule and fluoride ion. However, we compute the energy of the hydrogen bond with F- in a manner similar to that used in the energy calculations for the cation-molecule complex, i.e.j with respect to the formamide molecule and fluoride anion. The interaction energy is as follows: for Li+-formamide complex -62.3 kcal mol-I, for F--formamide system -51.5 kcal mol-’ and for the system Li+-formamide-F-205.8 kcal mol-‘. The calculated force constants are listed in Table 3. It is seen from Table 3 that all the complexes show considerable changes in force constants of the three stretching vibrations of the CO, CH and NH groups. The cation slightly decreases the CO frequency and increases the CH frequency. On the other hand, the anion decreases the CO as well as NH and CH frequencies. TABLE 2 Results of the energy 4-31G
calculations on the formamide complexes with ions
Species
Total energy (a.u.)
Formamide Li+-formamide F--formsmide Li+-formamide-FFluoride ion Lithium cation
-168.681588 -176.015114 -268.011430 -275.490751 -99.247824 -7.233262
226 TABLE 3 Force constants for the stretching vibrations in mdyn a-’ Species
kc0
kcri
kN?x
Formamide Li+-formamide F--formamide
13.8 13.6 11.8
5.80 6.07 5.77
Li+-formamide-F-
12.2
6.05
8.15 8.15 4.99 8.07 1.72 7.81 3.98
(sym NH,) (sym NH,) (sym NH,) (NH,) (NH,) (NH,) (NHt)
The lower CO bond strength in the lithium-cation complex and fluorideanion complex is also reflected in the evaluated equilibrium bond distances Rco. The CO bond on all the complexes are slightly longer than that in the formamide. molecule. Conversely, the CN bonds are slightly shorter. It should be pointed out that Ottersen and Jensen [Xi] and Ottersen et al. [ 181 obtained different results for the formamide-water [ 151 and formamidewater [ 15) and formamideammonia [ 181 systems. In these papers the kc0 values increase upon formation of the hydrogen bond. This result is rather puzzling. For the NH bond we calculated the force constants for the symmetric vibration and for the stretching one of the NH bond in the case of the complexes with anions. Their values are very different. It should be noted that both these quantities, the stretching force constants and the bond lengths, correspond only to the net effect of the rearrangement of electron density upon complex formation. However, it is interesting to look for the more detailed analysis of the electron density redistribution in the perturbed system. It can partly be achieved by means of an analysis of the atomic population. Table 4 presents the net atomic population for the investigated molecular systems. TABLE 4 The net atomic charges for foramide and its complexes Atom
Formamide
Lit-formamide
F--formamide
Li+-formamide-F-
C
0.587 -0.904 -0.612 0.170 0.373 0.386 0.0
0.700 -0.855 -0.831 0.261 0.428 0.428 0.869 0.123
0.565 -0.962 -0.740 0.133 0.521 0.309 0.825 -0.175
0.597 -0.868 -0.934 0.195 0.585 0.320 0.779 -0.676 -0.103
N 0 H Ht HC Li+ FTotal net charges in formamide
227
When comparing the net atomic populations in complexes with those in the parent formamide molecule, a considerable charge redistribution is noticed. An examination of the changes indicates that there is a charge transfer from anion to formamide molecule. In the complex with cations, however, electron density is transferred in the opposite direction. The increase in the polarity of the CN and CO region is observed: the oxygen atom gains electrons in all complexes, while the changes on the nitrogen and carbon atoms are in opposite directions upon the coordination of cation and anion. This picture of electron density changes in the formamide molecule, which results from the adduct formation, is more directly obtainable from the set of localized molecular orbit&. The localized MO are examined according to Boys [41]. In Table 5 we present the coordinates of the centroid of charge with respect to the molecular coordinate system for all the localized molecular orbit&. In the complex with a cation, the centre of charge for two CO LMO is shifted towards the oxygen atom, for two NH LMO towards the nitrogen atom. In the complex with an anion, one LMO which previously described the NH, bond is now identified, using critical level 0.080, as centered on the N atom. The complex OCH-NH2 - F- seems now to be the HOCNH-. . . H’. . . F- system. An interesting effect is observed when two ions are placed close to the parent formamide molecule. Now, one of the LMOs is centered on the N atom, but one of the LMOs which described the CO bond is centered only on the oxygen atom. It induces a change in the hybrydization of oxygen lone pairs, which now have the centre of charge located out of the molecular plane and the CO bond seems to be a single bond rather than a double one: Li’...HO---C--NH~o.H’.*.F-. Finally, the results presented above can be compared with some experimental observations in the Raman spectra. The model calculations refer to the l
l
TABLE 5 Centroid of charge (x, y, z in bohr) for formamide and its complexes LMO
formamide
Li+-formamide
F--formamide
Li+-formamide--F
ls0 ISN ISC NH+ NH; co co CN CN CH
1.89 0.00-1.31 0.00 0.00 0.00 0.00 0.00 0.00 -0.94 0.00 3.13 0.95 0.00 3.09 1.23 0.46-0.85 1.23 -0.46-0.85 0.00 -0.50 1.87 0.00 0.50 1.87 -1.28 0.00 -a.54 2.38 0.00 -0.99 1.75 0.00 -1.88 -
1.94 0.00 -1.32 0.00 0.00 2.48 0.00 0.00 0.00 -0.92 0.00 3.05 0.93 0.00 3.03 1.39 -0.44 -0.94 1.39 0.44 -0.94 0.00 0.51 1.75 0.00 -0.51 1.75 -1.20 0.00 -0.57 2.50 0.00 -1.03 1.85 0.00 -1.96 4.36 0.00 -3.44 -
1.87 0.00 -1.42 0.00 0.00 2.49 0.00 0.00 0.00 -0.81 0.00 3.0W 1.01 0.00 3.02 1.27-0.45~.97 1.27 0.45 -0.97 0.01-0.51 1.75 0.01 0.51 1.75 -1.30 0.00 -0.54 2.38 0.00-1.13 1.71 0.00-2.01 -4.17 0.00 4.75 -3.58 0.00 4.45 -4.47 -0.41 4.57 4.41 0.45 4.63 4.09 0.00 5.27
1.97 O.OO-1.44 0.00 0.00 2.41 0.00 0.00 0.00 -0.69 0.00 2.83a 1.01 0.00 2.93 1.18 0.04 -0.88 2.07-0.53-l.32b 0.02 0.53 1.53 0.02 AI.53 1.53 -1.22 0.00-0.61 2.39 0.37 -1.16 1.82 0.12 -2.07 4.49 0.00-3.17 4.07 0.00 4.25 -3.37 0.00 4.00 4.16 0.46 4.49 4.38 4.06 3.83 4.13 -0.40 4.57
IPO lP0
1sLi ISF 2sF 2pF 2pF 2pF
aLMO is centered on the N atom. bLMO is centered on the 0 atom.
228
isolated molecule and its isolated complex with cations and anions. These data should be compared rather with the gas-phase spectra of formamide. The only experimental data available at that moment are for the liquid solvent and the corresponding solutions of electrolytes. However, we can try to use the information following from the comparison of the gaseous and pure liquid formamide. Amide I band frequency It is proposed that, in the liquid phase, the formamide molecules tend to form, most probably, antiparallel, cyclic dimers [3]. It was found that the CO/amide I band frequency is 95 cm -’ lower in pure liquid formamide than in the gas phase [ 351. This latter observation suggests that the breaking of hydrogen bonds between self-associated formamide molecules in gas phase produces the amide I frequency increase. This is consistent with ab initio SCF calculations for hydrogen bonded cyclic dimer [ 121. When electrolytes are dissolved in formamide the frequencies of the amide I band are higher than in the pure liquid formamide [ 4,6]. However, in the ternary solutions, formamide + acetonitrile + salt, [6] for various salts the results show a decrease in the amide I band frequency. According to the authors [6, 391, these results prove that cations decrease the amide I frequency being linked to the CO group, while the anions increase the frequency changes provoked by cation. This conclusion comes from the assumption that in the binary solutions, formamide + acetonitrile, some of formamide dimers are broken by the formamide-acetonitrile interaction and, next, upon the dissolution of the electrolyte, the preferential solvation of ions by the amide can be found [3]. Theoretical results for the model systems presented in Table 3 show the small decrease of the CO force constant after the coordination of the cation through the CO bond and the rather drastic decrease after the interaction with fluoride anion. It should be kept in mind, however, that the CO stretching contributes only 65% to the amide I mode [36] or even less [38]. The rest is CN stretching vibration, 6 NH, and VN-n, which also changes after the coordination with a cation or anion. In spite of the approximation, assuming that the CO bond of the formamide molecule vibrates independently, there is a qualitative agreement between the computed frequency changes and those found in the spectra [6]. The CH stretching band The CH stretching band is influenced by ions in a different manner. As follows from Table 3, a cation causes an increase in its frequency but an anion decreases it. Nevertheless, there are no other calculations on the CH stretching frequency for the model system. Thus, we can compare our results obtained for the isolated system with the experimental data for solutions of
229
electrolytes in HCOND2. Raman spectra of the binary and ternary solutions of electrolytes indicate an increase of the CH stretching band caused by cations and a decrease caused by anions. Because the Yc_n stretching vibration is very characteristic [36], we can say, that these findings are in agreement with the calculations. The NH stretching band Both ions cause a decrease of the NH stretching band for the isolated system. The experimental results [4] indicate the lowering of the ND stretching frequency of the HCONHD upon dissolved the salts. Because both vibrations are characteristic [ 37, 381, also a lowering VN-HS~~. andv~--H~~~. of the symmetric NH2 stretching frequency is expected, in agreement with the presented calculations. ACKNOWLEDGEMENTS
We are indebted to Professor Z. Kecki for his interest and helpful comments on the manuscript. This work was supported, in part, by the Polish Academy of Sciences in connection with project MR.I.9. REFERENCES 1 D. E. Iresh and M. H. Brooker, in R. J. H. Clarke and R. E. Hester (Eds.), Advances in Infrared and Raman Spectroscopy, Vol. 2, Heyden, London, 19’76. 2 M. H. Baron and C. de Loze, J. Chim. Phys. Phys. Chim. Biol., 69 (1972) 1084. 3 M. H. Baron, H. Jaeschke, R. M. Moravie, C. de Loze and J. Corset, in Metal-Ligand Interactions in Organic Chemistry and Biochemistry, Part 1, D. Reidel, Dordrecht, 1977, p. 171. 4 J. Bukowska, Chem. Phys. Lett., 57 (1978) 624. 5 D. J. Gardiner, R. B. Girling and R. E. Hester, J. Chem. Sot. Faraday Trans. 2, 71 (1975) 709. 6 J. Bukowska and K. Miaskiewicz, J. Mol. Struct., 74 (1981) 1. 7 M. Dreyfus and A. Pullman, Theor. Chim. Acta, 19 (1970) 20. 8 A. Pulhnan, H. Berthold, C. Giessner-Prettre, J. F. Hinton and R. D. Harpool, J. Am. Chem. Sot., 100 (1978) 3991. 9 J. F. Hinton and R. D. Harpool, J. Am. Chem. Sot., 99 (1977) 349. 10 S. Yamabe, K. Kitaura and K. Nishimoto, Theor. Chim. Acta, 47 (1978) 111. 11 T. Ottersen and H. H. Jensen, J. Mol. Struct., 26 (1975) 355. 12 T. Ottersen, J. Mol. Struct., 26 (1975) 365. 13 F. A. Momany, R. F. McGuire, J. F. Yan and H. A. Scheraga, J. Phys. Chem., 74 (1970) 2424. 14 D. Peters and J. Peters, Int. J. Quantum Chem., 19 (1981) 1121. 15 T. Ottersen and H. H. Jensen, J. Mol. Struct., 26 (1975) 375; 28 (1975) 220. 16 J. E. Del Bene, J. Chem. Phys., 62 (1975) 1314,1961. 17 J. E. Del Bene, J. Am. Chem. Sot., 100 (1978) 1387,1395. 18 T. Ottersen, H. H. Jensen, R. Johansen and E. Wisloff-Nilssen, J. Mol. Struct., 30 (19760 379. 19 P. Otto, S. Suhai and J. Ladik, Int. J. Quantum Chem. Biol. Symp., 4 (1977) 451. 20 R. Janoschek, Theor. Chim. Acta, 32 (1973) 49.
230 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41
R. Janoschek, The Jerusalem Symposia, 6 (1974) 284. B. M. Rode and H. W. Preuss, Theor. Chim. Acta, 35 (1974) 369. R. Fusseneger and B. M. Rode, Chem. Phys. Lett., 44 (1976) 95. B. M. Rode and,R. Fusseneger, J. Chem. Sot. Faraday Trans. 2,71 (1975) 1958. D. N. Fuchs and B. M. Rode, Chem. Phys. Lett., 82 (1981) 517. B. M. Rode, in Metal-Ligand Interactions in Organic Chemistry and Biochemistry, Part 1, B. Reidel, Dordrecht, 1977, p. 127. P. Schuster, in L. Kevan and B. Webster (Eds.), Electron-Solvent and Anion-Solvent Interactions, Elsevier, Amsterdam, 1976, p. 259. J. Emdey, D. J. Jones, J. M. Miller, R. E. Overill and R. A. Waddilove, J. Am. Chem. Sot., 103 (1981) 24. R. Ditchfield, W. J. Hehre and J. A. Pople, J. Chem. Phys., 54 (1971) 724. P. C. Hariharan and J. A. Pople, Mol. Phys., 27 (1974) 209. M. Kitano and K. Kuchitsu, Bull. Chem. Sot. Jpn, 47 (1974) 67. P. Shuster, W. Jakubtz, W. Marius, Top. Current Chem., 60 (1975) 1. H. G. Hertz, H. Weingartner and B. M. Rode, Ber. Bunsenges. Phys. Chem., 79 (1975) 1190. R. D. Green, Can. J. Chem., 47 (1969) 2497. S. T. King, J. Phys. Chem., 75 (1971) 405. I. Suzuki, Bull. Chem. Sot. Jpn, 33‘(1969) 1359. Y. Sugawara, Y. Hamada, A. Y. Hirakawa, M. Tsuboi, S. Kato and K. Morokuma, Chem. Phys., 50 (1980) 105. E. D. Schmid, E. Brodbeck, J. Chem. Phys., 78 (1983) 1117. J. Bukowska, J. Mol. Struct., 98 (1983) 1. E. Hirota, R. Sugisaki, C. J. Nielsen and G. 0. Sorensen, J. Mol. Spectrosc., 49 (1974) 251. S. R. Boys, Rev. Mod. Phys., 32 (1960) 306.