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Hydrogen-bonding interaction between acetic acid and pyridine 1:1 complex
Journal of Molecular Structure 660 (2003) 113–118 www.elsevier.com/locate/molstruc
Hydrogen-bonding interaction between acetic acid and pyridine 1:1 complex Kemei Pei, Yimin Li, Haiyang Li* Laboratory of Environment Spectroscopy, Anhui Institute of Optics and Fine Mechanics, Chinese Academy of Sciences, P.O. Box 1125, Hefei 230031, People’s Republic of China Received 10 July 2003; revised 8 August 2003; accepted 12 August 2003
Abstract Acetic acid (CH3COOH) and pyridine (C5H5N) 1:1 complex was subjected to Density Functional Theory (DFT) calculations using 6-311G(d,p) basis set. Five equilibrium isomers were located, and the most stable structure has a strong N· · ·H – O hydrogen-bonding interaction. A normal-mode analysis of the vibrations of the five isomers was carried out and compared with the experimental values. The calculations indicate that solvents enhance significantly the strength of hydrogen bond as shown by the decrease of the N – H distance and appreciable red shift of the H– O vibration mode and blue shift of N· · ·H – O vibration mode. q 2003 Elsevier B.V. All rights reserved.
1. Introduction The molecules that constitute a cluster are usually bound to one another via weak van der Waals interactions or stronger ones such as hydrogen bonds. Researching into hydrogen-bonded molecular clusters is crucial with a view to understanding a wide variety of chemical and biochemical processes. Hydrogen bonds not only dominate the influence of the aqueous environment on molecular structure, but also play a key role in determining the structure and the function of biomolecules, such as proteins and nucleic acids [1 –5]. Acid – pyridine systems are the typical hydrogenbonding systems existing two different proton* Corresponding author. Fax: þ 86-551-559-1550. E-mail address: [email protected] (H. Li). 0022-2860/$ - see front matter q 2003 Elsevier B.V. All rights reserved. doi:10.1016/j.molstruc.2003.08.010
limiting structures, namely, OH· · ·N , O2 · · ·Hþ N; which yield hydrogen-bonding and ionic interpolymer interactions, respectively [6 – 15]. Acetic acid and pyridine complex serves as a model system for larger aggregates, and the studies of it are mostly concentrated on the vibration spectroscopy and solvent effect on experimental level. The IR bands of vibrations involving proton movement depend on the dielectric constant largely. Though the existence of such dependence seems to be a well-known fact, the number of works, where direct comparative study of the same complex in gaseous phase and solutions of different dielectric constant were performed, are not numerous. Most part of these data concern the systems with weak and middle strength H-bonds, the transition of such a complex from gas phase to solution, as well as an increase of
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solvent polarity, leads usually to an additional frequency shift of some characterized frequencies, such as O – H, CyO [7 – 14]. The experimental findings indicate that the basis of the interaction between acids and pyridine is not well understood. Therefore, theoretical studies are required to shed light on this interesting problem. In this paper, Density Function Theory (DFT) will be used to study the equilibrium structures and the solvent effect of the most stable structure of acetic acid and pyridine 1:1 complex.
2. Computational methods All calculations were carried out with the 98 package [16]. It has already been proven that the DFT method supplies excellent results in most hydrogen-bonding systems [17 – 20]. The Becke proposed hybrid (B3) together with the LYP correlation functional and the basis set 6311G(d, p) were chosen for this work [21 – 22]. Frequencies were calculated at this level of theory and the corresponding zero-point energy (ZPE) corrections were made to the total energy. All calculated frequencies have been scaled by 0.96 as recommended by Wong [23]. The binding energies were evaluated with the ZPE corrected energies as De ¼ ðEA þ EB Þ 2 EAB ; where AB stands for the complex, and A and B for each of the two monomer molecules of the complex. For more accurate understanding the nature of the hydrogen bond, the binding energies were corrected by the Basis Set Superposition Error (BSSE), which was estimated using the counterpoise method (CP) [24]. The Natural Bond Orbital (NBO) analysis was used to understand better the nature of the corresponding intermolecular interactions. NBO analysis was performed on the NBO program in the GAUSSIAN 98 package [25 –26]. The NBO program uses the polyatomic wave function and localizes the molecular orbitals, and supplies data that is in well agreement with the concepts of Lewis structures and basic Pauling – Slater – Coulson picture of bond hybridization and polarization [25,27]. GAUSSIAN
3. Results and discussion 3.1. Geometries Five equilibrium structures, which are shown in Fig. 1, were found on the potential surface of acetic acid and pyridine 1:1 complex. Structure I is the most stable structure characterized in this study. Inspection of its geometry reveals the presence of a strong hydrogen-bonding interaction, namely, N· · ·H –O, and the N – H distance is 0.1833 nm. Structure II is the second stable isomer shown in Fig. 1, its energy is 27.15 kJ/mol higher than I. There are two very weak hydrogen-bonding interactions, namely, N· · ·H – C and C –H· · ·O. The N – H distance in N· · ·H – C is 0.2414 nm, and the H –O distance in C– H· · ·O is 0.2370 nm. Additionally, the two hydrogen bonds are in a plane. Structure III is the rotation isomer of II, which forms from the acetic monomer rotating around the pyridine monomer plane. As well as isomer II, isomer III has two hydrogen bonds, namely, N· · ·H – C and C –H· · ·O. The N – H distance and the H –O distance are 0.2456 and 0.2440 nm, respectively, which is longer than the correspondent
Fig. 1. The equilibrium structures of CH3COOH and C5H5N complex isomers (the relative energies showed in ‘[ ]’).
K. Pei et al. / Journal of Molecular Structure 660 (2003) 113–118
distance in isomer II. The stability of isomer III is weaker than II, which has a relative 29.32 kJ/mol energy to isomer I. Structure IV and V are formed from the two O atoms in acetic acid combining with the two H atoms of pyridine, namely, the two C – H· · ·O hydrogen bonds interaction. The different position of the two C – H· · ·O bonds leads to the different stability. The H – O distances of the two hydrogen bonds in IV are 0.2769 and 0.2460 nm, and the corresponding distances of V are 0.2716 and 0.2490 nm. The stability of these two isomers is much weaker than that of structure I– III, the energies of IV and V are 32.79 and 33.52 kJ/mol higher than the most stable structure I, respectively. 3.2. Energies As shown in Table 1, the bonding energies De of the five isomers is: 42.52 (I), 17.08 (II), 14.87 (III), 12.69 (IV) and 12.02 (V) kJ/mol. After the BSSE is: 32.71(I), correcting, the bonding energies Dcorr e 9.16(II), 7.86(III), 4.85(IV) and 4.43(V) kJ/mol. Comparing De with Dcorr e ; we can conclude that the BSSE value is large enough and should be considered. For example, the BSSE correction scale of V is about 170%. The bonding energies of I is far larger than that of the other four structures, namely, the interaction between the two monomers of I is far stronger than those of the other four structures.
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Analysis of the NBO second-order interaction energies DEijð2Þ of these hydrogen-bonding interactions points to a different interpretation. DEijð2Þ is calculated as DEijð2Þ ¼ 2
^ fj ll2 lkfi lFl 1i 2 1j
where 1i and 1j correspond to the energy eigenvalues of the donor molecular orbital fi and the acceptor molecular orbital fj ; DEijð2Þ measures the strength of the donor – acceptor interaction between orbitals fi and fj and appears to be well suited to examine hydrogen-bonding interactions [28]. Inspection of Table 1 reveals that the O· · ·H – N hydrogen bond of I is remarkably strong, with a large second-order interaction energy of 78.41 kJ/mol. There are two weak hydrogen bond in structure II, namely, N· · ·H – C and O· · ·H – C, and their corresponding second-order interaction energies are 9.96 and 4.44 kJ/mol, respectively. Obviously, the strength of N· · ·H – C is larger than that of the O· · ·H – C. Therefore, the bulk of the bonding interaction in II should be ascribed to the N· · ·H –C hydrogen-bonding interaction, the contribution of O· · ·H – C hydrogenbonding interaction is considerably smaller. Similar to II, III has two weak hydrogen bonds too, and the second-order interaction energies are 8.70 kJ/mol (N· · ·H –C) and 1.72 kJ/mol (O· · ·H – C), the N· · ·H – C hydrogen bond plays a key role in III. Their corresponding second-order interaction energies are
Table 1 ð2Þ Binding energies De =Dcorr e ; donor orbitals fi ; acceptor orbitals fj and their corresponding second-order interaction energies DEij of the hydrogen-bonded structures DEijð2Þ ðkJ=molÞ
Isomers
De (kJ mol21)
Dcorr ðkJ mol21 Þ e
fi ! fj
I
42.52
32.71
n N ! dp O 2 H p
78.41
II
17.08
9.16
nN!d C2H n O ! dp C 2 H
9.96 4.44
III
14.87
7.86
n N ! dp C 2 H n O ! dp C 2 H
8.70 1.72
IV
12.69
4.85
n O ! dp C 2 H n O ! dp C 2 H
2.97 1.00
V
12.02
4.43
n O ! dp C 2 H n O ! dp C 2 H
2.30 1.30
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smaller than that of II, therefore, the interaction strength of the monomers in II is stronger than that in III. Weakly bond clusters IV and V, with binding energies of less than 5 kJ/mol, as shown in Fig. 1 and Table 1, has accordingly small second-order interaction energies.
acid and pyridine complex is formed. The agreement between our calculation and the experiment is quite well [8]. From Table 2, the largest red shift of CyO stretching vibrational mode is also in structure I. The calculated value of DnCyO is 2 26.1 cm21, which is also agree well with the observed value 2 20 ^ 10 cm21 [8].
3.3. Vibration spectra 3.4. Solvent effect Probing the structure of the acid –pyridine systems has turned out to be rather challenging, and the familiar method is to study their vibration spectra. In this paper, three characterized vibration mode frequencies nfull-sym (the full symmetrical stretching mode of pyridine), nO – H (the stretching mode of O – H) and nCyO (the stretching mode of CyO) have been discussed, and the corresponding frequency shifts are labeled as: Dnfull-sym ; DnO – H and DnCyO : As shown in Table 2, the Dnfull-sym of structure I is 7.7 cm21, which is the largest in the five structures. Isomer I is the most stable cluster, therefore, it has the largest abundance in acetic and pyridine 1:1 complex, and the blue shift phenomenon would occur near the fully symmetrical stretching bands of pyridine. Table 2 reveals that a large red shift would occur on the stretching vibration mode characterized spectra line of O –H in Structure I, the DnO – H is 2 543 cm21, the other four structures have much smaller DnO – H relative to I. The N· · ·H – O is a very strong hydrogen bond, and its formation can lead to the large frequency shift of the O –H stretching vibration frequency. Therefore, the existing of I can be examined based on the large red shift of nO – H : The experimental determined gas phase frequencies of nO – H in acetic acid and the complex are 3580, 3000 cm21 (peak maximum), so the frequency shift DnO – H is about 2 580 ^ 50 cm21 when 1:1 acetic
The solvent effect of acetic acid and pyridine 1:1 complex were explored by the Onsager reaction model in Self-Consistent Reaction Field (SCRF) method, and we concentrate on the vibration spectra, charge distribution and the hydrogen bond geometry parameters of the most stable structure I. Fig. 2 collects the shift frequency value of the characterized vibration mode in different dielectric constant solvent. As shown on Fig. 2, the nO – H frequency has a strong dependence on the dielectric constant 1; and the absolute value of its frequency shift value DnO – H is far larger than those of the other three ðDnCyO ; Dnfull-sym ; DnOH· · ·N Þ: When 1 is within 1.0 – 15.0, nO – H decreases sharply, but when 1 exceeds 15.0, nO – H decreases very slowly, and reach almost constant value about 2 370 cm21. Our calculations also predict a tiny red shift of nCyO with the dielectric constant 1 increasing. Similar to nO – H ; the frequency of CyO stretching decreases obviously when 1 is in 1.0– 15.0, and weak solvent effect after 1
Table 2 Frequency shifts of the characterized vibration modes Isomers
Dnfull-sym
DnO – H
DnCyO
I II III IV V
7.7 20.4 1.1 2.0 5.6
2443.1 2.2 11.1 22.1 22.5
226.1 215.7 224.7 210.3 28.4
Fig. 2. Characterized frequencies depended on the dielectric constant of solvent.
K. Pei et al. / Journal of Molecular Structure 660 (2003) 113–118
exceeding 15.0. The nfull-sym of pyridine monomer in I shows very weak solvent effect, and it increases slowly with 1 increasing. Experimental studies of the O – H vibration band of the carboxyl indicate that it has a strong dependence on 1 as reported by Drichko et al. and they also found that the transition from the gas phase to solution decreases the frequency of the O – H vibration band, for example, which is in good agreement with our calculations [8]. Drichko et al. measured the first moment M1 of the nOH· · ·N band in a number of solvents of different polarity, the M1 values change from hexane ð1 ¼ 1:89Þ 2672 cm21, carbon tetrachloride ð1 ¼ 2:23Þ 2577 cm21, hexaclorobutadiene ð1 ¼ 2:6Þ 2558 cm 21, butyl chloride ð1 ¼ 7:39Þ 2434 cm21, the frequency shift from hexane to butyl chloride is about 228 ^ 50 cm21; our calculated shift is 195 cm21, which agrees very well with the observed. Our calculated red shift of nCyO is and is comparable with observed exper13 cm21 , imental shift nmax 15 ^ 5 cm21. The nOH· · ·N increases slowly with 1 increasing, as shown in Fig. 2, which is also agree with the experiment results. Such a frequency shift indicates that the strength of the O – H· · ·N hydrogen bond interaction increases with the dielectric constant of the solvent. In order to understand the solvent effect of the hydrogen in structure 1, we also have calculated the charge distribution and geometry parameters of it in solution of different dielectric constant, as shown inTable 3. Negative charge on N and O of O – H· · ·N hydrogen bond increase and positive charge on H atom increases, when 1 increases. It is well known that more large polarity of the chemical bond is, and it will be more stable in the solvent with more large polarity. Our calculations illustrate the point. It is clear from the inspection of bond lengths given in Table 3, that solvent existing leads to decrease substantially the N –H distance and increase the O – H distance of O – H· · ·N hydrogen bond. We considered that just for shortening the N –H distance and lengthening the O – H distance lead to a strengthening of the hydrogen bond interaction, and this is best seen by inspection of appreciable red shift of the H – O vibration mode and blue shift of N· · ·H – O vibration mode from the above discussion.
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Table 3 Charge distribution and bond length of hydrogen-bond 1
N
H
O
N–H
O–H
Gas 2.0 3.0 5.0 7.0 9.0 11.0 13.0 15.0 17.0 19.0 21.0 23.0 27.0 30.0
20.407 20.420 20.426 20.432 20.435 20.436 20.437 20.438 20.439 20.439 20.439 20.439 20.440 20.440 20.441
0.282 0.289 0.292 0.295 0.296 0.297 0.298 0.298 0.298 0.299 0.299 0.299 0.299 0.299 0.299
20.325 20.326 20.328 20.329 20.330 20.330 20.330 20.330 20.331 20.331 20.331 20.331 20.331 20.331 20.331
1.830 1.790 1.770 1.750 1.740 1.740 1.730 1.730 1.730 1.730 1.730 1.730 1.730 1.720 1.720
0.987 0.993 0.997 1.001 1.002 1.004 1.004 1.005 1.006 1.006 1.006 1.006 1.006 1.007 1.007
4. Conclusion Acetic acid and pyridine 1:1 complex has five equilibrium structures predicted by the DFT theory. Strong N· · ·H – O hydrogen bond is formed in the most stable structure and the calculated bonding energy of it is 32.71 kJ/mol. Vibration spectra analysis predicts that shifts of the characterized vibration bands nfull-sym ; nO – H and nCyO ; which agree well with experimental observation. The calculation also indicates that solvent effect enhances N· · ·H – O hydrogen-bonding interaction, and leads to a large red shift of the H –O vibration mode, a small red shift of CyO vibration mode, and the blue shift of N· · ·H – O hydrogen-bonding vibration mode. Our DFT studies of the pyridine and acetic acid complex will help to assign the complicated experimental infrared spectra of the complex both in gas phase and in solution.
Acknowledgements This work was supported by the National Natural Science Foundation (Grant No. 20073042) and the Director Research Grant of Anhui Institute of Optics and Fine Mechanics (2003).
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References [1] G.A. Jeffrey, An introduction to Hydrogen Bonding, Oxford University Press, New York, 1997. [2] S. Scheiner, Hydrogen Bonding. A Theoretical Perspective, Oxford University Press, New York, 1997. [3] J.M. Ugalde, I. Alkorta, J. Elguero, Angrew. Chem. Int. Ed. Engl. 39 (2000) 717. [4] E.N. Baker, R.E. Hubbard, Prog. Biophys. Mol. Biol. 44 (1984) 97. [5] M. Jiang, X. Qun, W. Qin, L. Fei, Macromolecules 28 (1995) 710. [6] V. Berl, I. Huc, R.G. Khoury, M.J. Krische, J.M. Lehn, Nature 407 (2000) 720. [7] M.J. Fernandez-Berridi, J.J. Iruin, P. Iriondo, Macromolecules 29 (1996) 5605. [8] N.V. Drichko, G.Y. Kerenskaia, V.M. Schreiber, J. Mol. Struct. 477 (1999) 127. [9] M.J. Fernandez-Berridi, J.J. Iruin, L. Irusta, M.M. Jose, M.U. Jesus, J. Phys. Chem. A 106 (2002) 4187. [10] M.M. Coleman, P.C. Painter, Appl. Spectrosc. Rev. 20 (1984) 255. [11] R. Langner, G. Zundel, J. Chem. Soc. Faraday Trans. 91 (1995) 3831. [12] A.I. Kulbida, V.M. Schreiber, J. Mol. Struct. 47 (1978) 323. [13] I.V. Gerasimov, A.I. Kulbida, K.G. Tokhadze, V.M. Schreiber, Zhurn. Prikl. Spectr.(Russ.) 32 (1980) 1066. [14] V.A. Mikheev, V.M. Schreiber, Opt. i Spectr. 57 (1984) 3. [15] P. Hobza, H.L. Selzle, E.W. Schlag, Chem. Rev. 94 (1994) 1767. [16] M.J. Frisch, G.W. Trucks, H.B. Schlegel, G.E. Scuseria, M.A. Robb, J.R. Cheeseman, V.G. Zakrzewski, J.A. Montgomery,
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Hydrogen-bonding interaction between acetic acid and pyridine 1:1 complex Kemei Pei, Yimin Li, Haiyang Li* Laboratory of Environment Spectroscopy, Anhui Institute of Optics and Fine Mechanics, Chinese Academy of Sciences, P.O. Box 1125, Hefei 230031, People’s Republic of China Received 10 July 2003; revised 8 August 2003; accepted 12 August 2003
Abstract Acetic acid (CH3COOH) and pyridine (C5H5N) 1:1 complex was subjected to Density Functional Theory (DFT) calculations using 6-311G(d,p) basis set. Five equilibrium isomers were located, and the most stable structure has a strong N· · ·H – O hydrogen-bonding interaction. A normal-mode analysis of the vibrations of the five isomers was carried out and compared with the experimental values. The calculations indicate that solvents enhance significantly the strength of hydrogen bond as shown by the decrease of the N – H distance and appreciable red shift of the H– O vibration mode and blue shift of N· · ·H – O vibration mode. q 2003 Elsevier B.V. All rights reserved.
1. Introduction The molecules that constitute a cluster are usually bound to one another via weak van der Waals interactions or stronger ones such as hydrogen bonds. Researching into hydrogen-bonded molecular clusters is crucial with a view to understanding a wide variety of chemical and biochemical processes. Hydrogen bonds not only dominate the influence of the aqueous environment on molecular structure, but also play a key role in determining the structure and the function of biomolecules, such as proteins and nucleic acids [1 –5]. Acid – pyridine systems are the typical hydrogenbonding systems existing two different proton* Corresponding author. Fax: þ 86-551-559-1550. E-mail address: [email protected] (H. Li). 0022-2860/$ - see front matter q 2003 Elsevier B.V. All rights reserved. doi:10.1016/j.molstruc.2003.08.010
limiting structures, namely, OH· · ·N , O2 · · ·Hþ N; which yield hydrogen-bonding and ionic interpolymer interactions, respectively [6 – 15]. Acetic acid and pyridine complex serves as a model system for larger aggregates, and the studies of it are mostly concentrated on the vibration spectroscopy and solvent effect on experimental level. The IR bands of vibrations involving proton movement depend on the dielectric constant largely. Though the existence of such dependence seems to be a well-known fact, the number of works, where direct comparative study of the same complex in gaseous phase and solutions of different dielectric constant were performed, are not numerous. Most part of these data concern the systems with weak and middle strength H-bonds, the transition of such a complex from gas phase to solution, as well as an increase of
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K. Pei et al. / Journal of Molecular Structure 660 (2003) 113–118
solvent polarity, leads usually to an additional frequency shift of some characterized frequencies, such as O – H, CyO [7 – 14]. The experimental findings indicate that the basis of the interaction between acids and pyridine is not well understood. Therefore, theoretical studies are required to shed light on this interesting problem. In this paper, Density Function Theory (DFT) will be used to study the equilibrium structures and the solvent effect of the most stable structure of acetic acid and pyridine 1:1 complex.
2. Computational methods All calculations were carried out with the 98 package [16]. It has already been proven that the DFT method supplies excellent results in most hydrogen-bonding systems [17 – 20]. The Becke proposed hybrid (B3) together with the LYP correlation functional and the basis set 6311G(d, p) were chosen for this work [21 – 22]. Frequencies were calculated at this level of theory and the corresponding zero-point energy (ZPE) corrections were made to the total energy. All calculated frequencies have been scaled by 0.96 as recommended by Wong [23]. The binding energies were evaluated with the ZPE corrected energies as De ¼ ðEA þ EB Þ 2 EAB ; where AB stands for the complex, and A and B for each of the two monomer molecules of the complex. For more accurate understanding the nature of the hydrogen bond, the binding energies were corrected by the Basis Set Superposition Error (BSSE), which was estimated using the counterpoise method (CP) [24]. The Natural Bond Orbital (NBO) analysis was used to understand better the nature of the corresponding intermolecular interactions. NBO analysis was performed on the NBO program in the GAUSSIAN 98 package [25 –26]. The NBO program uses the polyatomic wave function and localizes the molecular orbitals, and supplies data that is in well agreement with the concepts of Lewis structures and basic Pauling – Slater – Coulson picture of bond hybridization and polarization [25,27]. GAUSSIAN
3. Results and discussion 3.1. Geometries Five equilibrium structures, which are shown in Fig. 1, were found on the potential surface of acetic acid and pyridine 1:1 complex. Structure I is the most stable structure characterized in this study. Inspection of its geometry reveals the presence of a strong hydrogen-bonding interaction, namely, N· · ·H –O, and the N – H distance is 0.1833 nm. Structure II is the second stable isomer shown in Fig. 1, its energy is 27.15 kJ/mol higher than I. There are two very weak hydrogen-bonding interactions, namely, N· · ·H – C and C –H· · ·O. The N – H distance in N· · ·H – C is 0.2414 nm, and the H –O distance in C– H· · ·O is 0.2370 nm. Additionally, the two hydrogen bonds are in a plane. Structure III is the rotation isomer of II, which forms from the acetic monomer rotating around the pyridine monomer plane. As well as isomer II, isomer III has two hydrogen bonds, namely, N· · ·H – C and C –H· · ·O. The N – H distance and the H –O distance are 0.2456 and 0.2440 nm, respectively, which is longer than the correspondent
Fig. 1. The equilibrium structures of CH3COOH and C5H5N complex isomers (the relative energies showed in ‘[ ]’).
K. Pei et al. / Journal of Molecular Structure 660 (2003) 113–118
distance in isomer II. The stability of isomer III is weaker than II, which has a relative 29.32 kJ/mol energy to isomer I. Structure IV and V are formed from the two O atoms in acetic acid combining with the two H atoms of pyridine, namely, the two C – H· · ·O hydrogen bonds interaction. The different position of the two C – H· · ·O bonds leads to the different stability. The H – O distances of the two hydrogen bonds in IV are 0.2769 and 0.2460 nm, and the corresponding distances of V are 0.2716 and 0.2490 nm. The stability of these two isomers is much weaker than that of structure I– III, the energies of IV and V are 32.79 and 33.52 kJ/mol higher than the most stable structure I, respectively. 3.2. Energies As shown in Table 1, the bonding energies De of the five isomers is: 42.52 (I), 17.08 (II), 14.87 (III), 12.69 (IV) and 12.02 (V) kJ/mol. After the BSSE is: 32.71(I), correcting, the bonding energies Dcorr e 9.16(II), 7.86(III), 4.85(IV) and 4.43(V) kJ/mol. Comparing De with Dcorr e ; we can conclude that the BSSE value is large enough and should be considered. For example, the BSSE correction scale of V is about 170%. The bonding energies of I is far larger than that of the other four structures, namely, the interaction between the two monomers of I is far stronger than those of the other four structures.
115
Analysis of the NBO second-order interaction energies DEijð2Þ of these hydrogen-bonding interactions points to a different interpretation. DEijð2Þ is calculated as DEijð2Þ ¼ 2
^ fj ll2 lkfi lFl 1i 2 1j
where 1i and 1j correspond to the energy eigenvalues of the donor molecular orbital fi and the acceptor molecular orbital fj ; DEijð2Þ measures the strength of the donor – acceptor interaction between orbitals fi and fj and appears to be well suited to examine hydrogen-bonding interactions [28]. Inspection of Table 1 reveals that the O· · ·H – N hydrogen bond of I is remarkably strong, with a large second-order interaction energy of 78.41 kJ/mol. There are two weak hydrogen bond in structure II, namely, N· · ·H – C and O· · ·H – C, and their corresponding second-order interaction energies are 9.96 and 4.44 kJ/mol, respectively. Obviously, the strength of N· · ·H – C is larger than that of the O· · ·H – C. Therefore, the bulk of the bonding interaction in II should be ascribed to the N· · ·H –C hydrogen-bonding interaction, the contribution of O· · ·H – C hydrogenbonding interaction is considerably smaller. Similar to II, III has two weak hydrogen bonds too, and the second-order interaction energies are 8.70 kJ/mol (N· · ·H –C) and 1.72 kJ/mol (O· · ·H – C), the N· · ·H – C hydrogen bond plays a key role in III. Their corresponding second-order interaction energies are
Table 1 ð2Þ Binding energies De =Dcorr e ; donor orbitals fi ; acceptor orbitals fj and their corresponding second-order interaction energies DEij of the hydrogen-bonded structures DEijð2Þ ðkJ=molÞ
Isomers
De (kJ mol21)
Dcorr ðkJ mol21 Þ e
fi ! fj
I
42.52
32.71
n N ! dp O 2 H p
78.41
II
17.08
9.16
nN!d C2H n O ! dp C 2 H
9.96 4.44
III
14.87
7.86
n N ! dp C 2 H n O ! dp C 2 H
8.70 1.72
IV
12.69
4.85
n O ! dp C 2 H n O ! dp C 2 H
2.97 1.00
V
12.02
4.43
n O ! dp C 2 H n O ! dp C 2 H
2.30 1.30
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smaller than that of II, therefore, the interaction strength of the monomers in II is stronger than that in III. Weakly bond clusters IV and V, with binding energies of less than 5 kJ/mol, as shown in Fig. 1 and Table 1, has accordingly small second-order interaction energies.
acid and pyridine complex is formed. The agreement between our calculation and the experiment is quite well [8]. From Table 2, the largest red shift of CyO stretching vibrational mode is also in structure I. The calculated value of DnCyO is 2 26.1 cm21, which is also agree well with the observed value 2 20 ^ 10 cm21 [8].
3.3. Vibration spectra 3.4. Solvent effect Probing the structure of the acid –pyridine systems has turned out to be rather challenging, and the familiar method is to study their vibration spectra. In this paper, three characterized vibration mode frequencies nfull-sym (the full symmetrical stretching mode of pyridine), nO – H (the stretching mode of O – H) and nCyO (the stretching mode of CyO) have been discussed, and the corresponding frequency shifts are labeled as: Dnfull-sym ; DnO – H and DnCyO : As shown in Table 2, the Dnfull-sym of structure I is 7.7 cm21, which is the largest in the five structures. Isomer I is the most stable cluster, therefore, it has the largest abundance in acetic and pyridine 1:1 complex, and the blue shift phenomenon would occur near the fully symmetrical stretching bands of pyridine. Table 2 reveals that a large red shift would occur on the stretching vibration mode characterized spectra line of O –H in Structure I, the DnO – H is 2 543 cm21, the other four structures have much smaller DnO – H relative to I. The N· · ·H – O is a very strong hydrogen bond, and its formation can lead to the large frequency shift of the O –H stretching vibration frequency. Therefore, the existing of I can be examined based on the large red shift of nO – H : The experimental determined gas phase frequencies of nO – H in acetic acid and the complex are 3580, 3000 cm21 (peak maximum), so the frequency shift DnO – H is about 2 580 ^ 50 cm21 when 1:1 acetic
The solvent effect of acetic acid and pyridine 1:1 complex were explored by the Onsager reaction model in Self-Consistent Reaction Field (SCRF) method, and we concentrate on the vibration spectra, charge distribution and the hydrogen bond geometry parameters of the most stable structure I. Fig. 2 collects the shift frequency value of the characterized vibration mode in different dielectric constant solvent. As shown on Fig. 2, the nO – H frequency has a strong dependence on the dielectric constant 1; and the absolute value of its frequency shift value DnO – H is far larger than those of the other three ðDnCyO ; Dnfull-sym ; DnOH· · ·N Þ: When 1 is within 1.0 – 15.0, nO – H decreases sharply, but when 1 exceeds 15.0, nO – H decreases very slowly, and reach almost constant value about 2 370 cm21. Our calculations also predict a tiny red shift of nCyO with the dielectric constant 1 increasing. Similar to nO – H ; the frequency of CyO stretching decreases obviously when 1 is in 1.0– 15.0, and weak solvent effect after 1
Table 2 Frequency shifts of the characterized vibration modes Isomers
Dnfull-sym
DnO – H
DnCyO
I II III IV V
7.7 20.4 1.1 2.0 5.6
2443.1 2.2 11.1 22.1 22.5
226.1 215.7 224.7 210.3 28.4
Fig. 2. Characterized frequencies depended on the dielectric constant of solvent.
K. Pei et al. / Journal of Molecular Structure 660 (2003) 113–118
exceeding 15.0. The nfull-sym of pyridine monomer in I shows very weak solvent effect, and it increases slowly with 1 increasing. Experimental studies of the O – H vibration band of the carboxyl indicate that it has a strong dependence on 1 as reported by Drichko et al. and they also found that the transition from the gas phase to solution decreases the frequency of the O – H vibration band, for example, which is in good agreement with our calculations [8]. Drichko et al. measured the first moment M1 of the nOH· · ·N band in a number of solvents of different polarity, the M1 values change from hexane ð1 ¼ 1:89Þ 2672 cm21, carbon tetrachloride ð1 ¼ 2:23Þ 2577 cm21, hexaclorobutadiene ð1 ¼ 2:6Þ 2558 cm 21, butyl chloride ð1 ¼ 7:39Þ 2434 cm21, the frequency shift from hexane to butyl chloride is about 228 ^ 50 cm21; our calculated shift is 195 cm21, which agrees very well with the observed. Our calculated red shift of nCyO is and is comparable with observed exper13 cm21 , imental shift nmax 15 ^ 5 cm21. The nOH· · ·N increases slowly with 1 increasing, as shown in Fig. 2, which is also agree with the experiment results. Such a frequency shift indicates that the strength of the O – H· · ·N hydrogen bond interaction increases with the dielectric constant of the solvent. In order to understand the solvent effect of the hydrogen in structure 1, we also have calculated the charge distribution and geometry parameters of it in solution of different dielectric constant, as shown inTable 3. Negative charge on N and O of O – H· · ·N hydrogen bond increase and positive charge on H atom increases, when 1 increases. It is well known that more large polarity of the chemical bond is, and it will be more stable in the solvent with more large polarity. Our calculations illustrate the point. It is clear from the inspection of bond lengths given in Table 3, that solvent existing leads to decrease substantially the N –H distance and increase the O – H distance of O – H· · ·N hydrogen bond. We considered that just for shortening the N –H distance and lengthening the O – H distance lead to a strengthening of the hydrogen bond interaction, and this is best seen by inspection of appreciable red shift of the H – O vibration mode and blue shift of N· · ·H – O vibration mode from the above discussion.
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Table 3 Charge distribution and bond length of hydrogen-bond 1
N
H
O
N–H
O–H
Gas 2.0 3.0 5.0 7.0 9.0 11.0 13.0 15.0 17.0 19.0 21.0 23.0 27.0 30.0
20.407 20.420 20.426 20.432 20.435 20.436 20.437 20.438 20.439 20.439 20.439 20.439 20.440 20.440 20.441
0.282 0.289 0.292 0.295 0.296 0.297 0.298 0.298 0.298 0.299 0.299 0.299 0.299 0.299 0.299
20.325 20.326 20.328 20.329 20.330 20.330 20.330 20.330 20.331 20.331 20.331 20.331 20.331 20.331 20.331
1.830 1.790 1.770 1.750 1.740 1.740 1.730 1.730 1.730 1.730 1.730 1.730 1.730 1.720 1.720
0.987 0.993 0.997 1.001 1.002 1.004 1.004 1.005 1.006 1.006 1.006 1.006 1.006 1.007 1.007
4. Conclusion Acetic acid and pyridine 1:1 complex has five equilibrium structures predicted by the DFT theory. Strong N· · ·H – O hydrogen bond is formed in the most stable structure and the calculated bonding energy of it is 32.71 kJ/mol. Vibration spectra analysis predicts that shifts of the characterized vibration bands nfull-sym ; nO – H and nCyO ; which agree well with experimental observation. The calculation also indicates that solvent effect enhances N· · ·H – O hydrogen-bonding interaction, and leads to a large red shift of the H –O vibration mode, a small red shift of CyO vibration mode, and the blue shift of N· · ·H – O hydrogen-bonding vibration mode. Our DFT studies of the pyridine and acetic acid complex will help to assign the complicated experimental infrared spectra of the complex both in gas phase and in solution.
Acknowledgements This work was supported by the National Natural Science Foundation (Grant No. 20073042) and the Director Research Grant of Anhui Institute of Optics and Fine Mechanics (2003).
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