INTERATION OF OXYGEN WITH TRANSITION METAL ENTERS - A DFT STUDY Hans Mikosch, a Ellie L. Uzunova b and Georgi St. Nikolov b a Institute of hemical Technologies and Analytics, Vienna University of Technology, Vienna 1060, AUSTRIA b Institute of General and Inorganic hemistry, Bulgarian Academy of Sciences, Sofia 1113, BULGARIA Molecular oxygen in its 3 Σ g ground state is a stable molecule and though the oxidation reactions of H 2, O and organic substances are exothermic, without the mediation of either a catalyst, or an enzyme, they proceed very slowly. 1-4 The excitation of dioxygen to the more reactive singlet state 1 g can be achieved photochemically assisted by a sensitizer. A more common method to promote the reactivity of molecular oxygen is the catalytic route: transition metal centers, embedded in zeolites or in large organic matrices, are able to trap reversibly and activate molecular oxygen. Most of the heterogeneous catalytic oxidation reactions are known to proceed via the Mars-van Krevelen mechanism, 5 according to which the re-oxidation of the transition metal center is the rate-determining step and the metal-oxygen bond strength is a crucial factor for the catalyst's activity. The interaction of molecular oxygen with such transition metal centers leads to the formation of end on η 1 -complex (superoxo form), side on η 2 -complex (peroxo form) in which the O O bond is weakened, or η 0 - complex in which O 2 is dissociated to form a dioxide. The dioxide formation reaction may proceed via the following reaction scheme: O 2 + M O 2(adsorbed) M O 2 (adsorbed) M + 2 O 2 (adsorbed) M 2+ MOO M(O 2 ) MO 2 Transition metal oxide, peroxide and superoxide clusters of general composition MO n, (n=1-4, M = r, Mn, Fe, o) and transition metal cations at extraframework cation sites in zeolite matrices were studied by the DFT formalism with the purpose of eliciting their relative stability, bonding scheme and the ordering of their electronic states. Molecular electrostatic potential maps were used to estimate the nucleophilic character of the oxygen atoms in the different clusters. The thermodynamic stability with respect to the detachment of either atomic or molecular oxygen was also evaluated. All calculations were performed with the B3LYP method, as implemented in the Gaussian 98 package. The bond populations were examined by natural bond orbital (NBO) analysis. The peroxide and superoxide clusters can be regarded as intermediate species in the formation of the oxide clusters, which represent the global minima in the M + O 2 system (M=r, Fe, Mn, o). MnO 2 in its 4 B 1 ground state was found to be a bent cluster, in agreement with previous experimental and
theoretical studies; bent geometry and high spin 5 B 2 ground state is predicted for FeO 2. The bonding scheme of quasilinear oo 2 in its 6 A 1 ground state is similar to that of strictly linear structure, the o O π-bonds being highly delocalized and formed with major participation of the O 2p orbitals (> 75%); the share of the o 3d orbitals being less than 20%. The M O bond lengths are comparable in both the peroxides and superoxides, O O bond lengths in superoxides being invariably shorter, than in peroxides. With the exception of the quasilinear MnOO ( 8 A' state), the MOO bond angle varies in the (110-140 ) range, much less, than the OMO angle of the dioxides. The net L M charge transfer increases in the order Mn < Fe < o < Ni. Molecular electrostatic potential (MEP) maps, based on calculated B1LYP densities were used to estimate the nucleophilic character of the oxygen atoms in the different clusters. 6 The oxygen s nucleophilic nature is related to the generated negative electric field and forms the following series: peroxides > superoxides > dioxides. Figure 1 shows that the positive electric field around the metal ions in the dioxides decays faster for dioxides than for peroxides. In peroxides the positive MEP contour on the side of the cation is closely matched by the negative electrostatic potential contour on the side of the two oxygen atoms. Since the iron dioxide has the smallest OMO angle, a single continuous region of high electron density around both oxygen atoms is formed. In Mn, o and Ni dioxides, the highest electron density region is split into two parts and there is a space of positive MEP between the oxygen atoms. The area of the highest electron density in MnOO and NiOO is concentrated at the terminal oxygen atom, which bears a lower partial charge, as compared to the oxygen atom, which forms the M O bond. REFERENES: [1] B.O. West, Polyhedron, 1989, 3, 219. [2] W.D. Woggon, Top. urr. hem., 1996, 184, 40. [3] P. A. ox, Transition Metal Oxides, Oxford University Press, 1992. [4] L. I. Simandi, Dioxygen Activation and Homogeneous atalytic Oxidation, Elsevier, Amsterdam 1991. [5] P. Mars, D. W. van Krevelen, hem. Engn. Sci., 1954, 3, 41. [6] E. Uzunova, G. Nikolov, H. Mikosch, hemphyshem 2004, 5, 100-109
0.02-0.03-0.05-0.02 0 0.03-0.05-0.05-0.03-0.02-0.05-0.03-0.02 0.03 0.02-0.03-0.02 MnO 2 ; 4 B 1 FeO 2 ; 5 B 2 oo 2 ; 6 A 1, quasilinear NiO 2 ; 5 A 1-0.05 0 0.04 0.035 0.03 0.03-0.04-0 -0.06-0.035-0.06-0.085-0.08-0.03-0.06-0.06-0.04-0.03 Mn(O 2 ); 6 A 1 Fe(O 2 ); 5 A 1 o(o 2 ); 4 A 1 Ni(O 2 ); 3 B 1 0.03-0.03 0.03-0.03-0.06 0.03-0.03 0.03-0.05-0.03-0.06-0.06 MnOO; 6 A' FeOO; 3 A' ooo; 2 A' NiOO; 1 A' Figure 5. Molecular electrostatic potential maps (au) of MO 2, M(O 2 ) and MOO clusters, (M=Mn, Fe, o, Ni).
QUANTENHEMISHE MODELLREHNUNGEN AN RESVERATROL-FORMALDEHYD REAKTIONEN Ferenc Billes 1,4, Ildikó Mohammed-Ziegler 2, Ernö Tyihák 3, Hans Mikosch 4 1 Budapester Universität für Technologie and Wirtschaftswissenschaft 2 hemisches Forschungscentrum, UAW, 3 Institut für Pflanzenschutz, UAW 4 Technische Universität Wien Resveratrol (3,5,4 -Trihydroxi-Trans-Stilben, Abb. 1.) ist eine biologisch sehr aktive Substanz. Darum wird sie medizinisch, biologisch und chemisch ausführlich untersucht. Sie kommt in Glukosidform in Weintrauben, in Erdnüssen, in vielen Heilpflanzen und im Rotwein vor. H 29 O 22 H 21 H 10 9 8 11 12 H H 23 24 7 20 14 H 25 13 H 15 16 1 2 3 H26 H O 19 6 5 4 H 28 O 18 H 17 H 27 Abb.1. Resveratrol Resveratrol wirkt - als Antioxidant in Organismen, - vermindernd die Konzentration des hochdichten holesterol, - beeinflusst Mitoptosis und Apoptosis; daher ist es anticarcinogen und deshalb auch chemisch schützend, - Formaldehyd adsorbierend in biologischen Zellen. Die zuletzt angeführte Wirkungsweise ist die Ursache für die gegenständlichen Untersuchungen der möglichen Reaktionen des Resveratrols mit Formaldehyd. Diese Arbeit beruht auf unserem Vortrag bei einer internationalen Konferenz [1] Wir verwendeten dazu auch unsere früheren Ergebnisse mit Pinosylvin (3,5-Dihydroxi-Trans-Stilben) [2]. Die Untersuchungen führten über zwei verschiedene Wege: schwingungsspektroskopische Messungen und quantenchemische Rechnungen. Die Infrarot-Spektren wurden mit einem Nicolet Magna 750 FT-IR, die Raman Spektren mit einem Nicolet FT-Raman Model 950 Spektrometer gemessen. Die quantenchemische Rechnungen wurden mit dem Gaussian 98 Programmpaket durchgeführt. Die Elektronendichtefunktional Theorie (DFT) wurde mit dem Becke3P86 Funktional und mit dem 6-31G* Basissatz angewendet. Für die Grundsubstanz Resveratrol wurde nach einer Geometrieoptimierung auch die Schwingungsfrequenzen und die entsprechenden Kraftkonstanten berechnet. Mit Hilfe der Kraftkonstanten und den für
Pinosylvin berechneten Skalierungsfaktoren wurden Spektren simuliert und die Symmetrie der einzelnen Grundschwingungen bestimmt. Die gemessenen und die simulierte Spektren sind in Abb. 2. dargestellt. Das für das isolierte Molekül berechnete Infrarot-Spektrum kann die intermolekulare Wechselwirkungen nicht wiedergeben. Darum entsteht ein grosser Unterschied zwischen den gemessenen und berechneten Spektren. Die Unterschiede zwischen den gemessenen und berechneten Intensitäten liegen einerseits in der verschiedenen Umgebung der Moleküle (kondensiert isoliert), andererseits in der Genauigkeit der Intensitätsrechnungen. 0.8 Measured infrared spectrum of resveratrol 12 Measured Raman spectrum of resveratrol 10 Absorbance 0.6 0.4 4 0.2 2 0.0 0 4000 3500 3000 2500 2000 1500 1000 500 0 Wavenumber / cm -1 4000 3500 3000 2500 2000 1500 1000 500 0 Wavenumber / cm-1 Intensity / a.u. 8 6 Intensity / a.u. 8 6 4 2 alculated infrared spectrum of resveratrol Intensity / A 4 amū 1 60 50 40 30 20 10 60 50 40 alculated Raman spectra of resveratrol perpendicular parallel polarisation 30 20 10 0 1800 1700 1600 1500 0 4000 3500 3000 2500 2000 1500 1000 500 0 Wavenumber / cm -1 0 4000 3500 3000 2500 2000 1500 1000 500 0 Wavenumber / cm -1 Abb.2. Gemessene und simulierte Resveratrol Schwingungsspektren Es wurden 7 Reaktionstypen zwischen Resveratrol und Formaldehyd angenommen. Resveratrol hat vier reaktionsfähige Gruppen (s. Abb. 1): drei Hydroxilgruppen und eine Vinylidengruppe zwischen den zwei Ringen. Formaldehyd kann in zwei Formen reagieren, als Formaldehyd selbst (Abb. 3) oder nach Wasseraddition als Formacetal (Abb. 4). Abb. 3 Abb. 4
Die gesamte elektronische Energie aller Ausgangsmoleküle (Resveratrol, Formaldehyd, Formacetal) und der Produktmoleküle (einzelne Produkte und Wassermolekül) wurde berechnet. Neben den isolierten Molekülen wurden die Rechnungen unter Einbeziehung der Wasserhülle wiederholt. Tabelle 1. zeigt die Reaktionen. Die letzte Rechnungen wurden mit Reaktionsfeld und unter Annahme der Isodensität und unveränderter Molekülstruktur durchgeführt. Tabelle 1. Die berechnete Reaktionen (optimierte Strukturen) Nummer Ausgang Produkt 1 + 2 + 3 + + + 4 + + 5 + + 6 + 7 +
Tabelle 1 zeigt, dass es um drei Reaktiongruppen geht, in welchen die Produkte Isomere sind. Die Reaktionsenergien wurden als die Differenz der Molekülenergien zwischen Produkt- und Ausgangsmoleküle berechnet. Tabelle 2. zeigt die berechneten Reaktionsenergien der Reaktionen zwischen isolierten Molekülen bzw. Molekülen in Wasser. Die Reaktionsenergien wurden in Lösung nur für eine Reaktion je Gruppe berechnet, um die Tendenz der Verschiebung erkennen zu können. Tabelle 2. Berechnete Reaktionsenergien Nummer 1 2 3 4 5 6 7 Bezeichnung der Reaktion Formacetal Addition an der Vinylidengruppe, 1. Formacetal Addition an der Vinylidengruppe, 2. Hydroxil Austausch für Aldehyd am 3 Hydroxil Austausch für Aldehyd am 5 Hydroxil Austausch für Aldehyd am 11 (4 ) Formaldehyd Addition an der Vinylidengruppe, 1. Formaldehyd Addition an der Vinylidengruppe, 2. Reaktionsenergie kj/mol isoliert in Wasser -73,45-70,49-10,41-45,47-46,00-50,20-16,99-45,81 18,97-46,56 Unter den Reaktionen isolierter Moleküle scheinen die Formacetaladditionen am besten stabilisierend. Dazu muss man aber bemerken, dass die Reaktionen zwischen isolierten Molekülen viel stärker exotherm zu sein scheinen, als dieselben in Lösung. Die Verschiebung in die endotherme Richtung kann so weit gehen, dass die Formaldehyd-Addition selbst endotherm zu werden scheint. Die Rechnungen werden mit der Vorausberechnung der Schwingungsfrequenzen der Produktmoleküle im isolierten Zustand als auch in Wasser fortgesetzt. Referenz 1. Billes, F., Mohammed-Ziegler, I., Tyihák, E., Mikosch, H.: Quantum chemical model calculations on the reactions of formaldehyde with resveratrol, 6th International onference on Role of Formaldehyde in Biological Systems, 12-16.10.2003, Pécs, Hungary. 2. Billes, F., Mohammed-Ziegler, I., Mikosch, H., Holmgren, A.: Vibrational spectroscopic and conformational analysis of pinosylvin. J. Phys. hem., 106, 6232-6241 (2002). 3. M.J. Frisch et al.: Gaussian 98, Revision A. 7., Gaussian, Inc., Pittsburgh, 1998.
Density Functional study of intramolecular hydrogen bonds in dicoumarols Natasha Trendafilova (Institute of General and Inorganic hemistry, Bulgarian Academy of Sciences, 1113 Sofia, Bulgaria Günther Bauer (Institute of hemical Technologies and Analytics, Technical University of Vienna, Vienna A-1060, Austria) 1. Introduction Both X-ray and NMR studies have shown that 3,3 -benzylidenebis(4- hydroxycoumarin) (PhD) and 3,3 -methylenebis(4-hydroxycoumarin) (D) are stabilized by two O-H O intramolecular HBs [1-2]. The formation of intramolecular HBs is an important factor in assisting the molecule to attain a suitable configuration for binding to an enzyme [2-4]. Some compounds with intramolecular HBs are both uncouplers and inhibitors of mitochondrial oxidative phosphorylation, while compounds, which can only form intermolecular HBs, are only uncouplers of oxidative phosphorylation. [5]. It was further found that large atomic charges give rise to an additional electrostatic stabilization of the system with strong H-bonds [6]. In the frame of the project we performed DFT calculations of PhD, D and the monomeric building-block, 4-hydroxycoumarin (4-H) at B3LYP/6-31G* level of theory, focusing on the type and the strength of the intramolecular O-H O HBs. The HB strength was estimated from different model structures with one or both HB ruptured. Natural population analysis data, the electron density (ρ b ) and Laplacian ( 2 ρ b ) properties as well as ν(o-h) red shifts were also used for estimation of the hydrogen-bonding interactions and the forces driving their formation. 2. Methods All calculations were performed using GAUSSIAN98 program package [7]. The reliably accurate description of weak interactions like HBs generally requires a treatment of electron correlation. Density functional calculations with Becke s three-parameter hybrid method using the correlation functional of Lee, Yang and Parr (B3LYP) have proved quite useful in this regard for studying system with HBs [8-10]. This method has been tested with different basis sets and has successfully been applied for estimation of
the relative strengths and preferred geometries of HBs in different systems [10-15]. Density functional theory offers an electron correlation correction frequently comparable to the second-order Møller-Plesset theory (MP2) or in certain cases, and for certain purposes even superior to MP2, but at considerably lower computational cost [10]. Due to the size of the systems studied the coast advantage that offers B3LYP method in comparison with MP2 was significant. Moreover, the lower computational cost of B3LYP as compared to other correlated methods allowed us to calculate the harmonic vibrational frequencies of the large systems studied. In order to obtain reasonable agreement between the calculated and experimental geometry parameters for the atoms, involved in the hydrogen bondings, different basis sets were tested in the course of the calculations: 6-31G*, 6-31+G** and 6-311G*. Analysis of the electronic charge density (ρ b ) and its Laplacian ( 2 ρ b ) was performed by means of the theory of molecular structure proposed by Bader and coworkers [16-19]. The calculated electron density, ρ b, and its second derivative, 2 ρ b, were used for describing the nature of the intramolecular O H-O bonds. The vibrational spectra of the compounds studied were calculated at the B3LYP/6-31G* optimized geometries. In order to assign the calculated frequencies to approximate vibrational descriptor, the vibrational modes have been analyzed by means of the atom movements, calculated in artesian coordinates. To improve the frequency shift estimation, scale factors were used: 0.956 for PhD and D and 0.972 for 4-H. 3. Results 3.1. Geometry parameters of PhD, D and 4H. The fully optimized molecular structures of PhD, D and 4-H, calculated at B3LYP/6-31G* level of theory, are shown in Fig. 1a, Fig. 1b and Fig. 1c, respectively. The calculated geometry parameters for PhD and D are in a reasonable agreement with available experimental data [1,20]. For PhD, two O-H O intramolecular HBs were found; each links a coumarin hydroxyl and carbonyl group (Fig. 1a). At B3LYP/6-31G* level both calculated O O distances,
a b c Fig. 1 Optimized geometries of 3,3 -benzylidenebis(4-hydroxycoumarin) (phenyldicoumarol, PhD), 3,3 -methylenebis(4-hydroxycoumarin) (dicoumarol, D) and the parent compound, 4- hydroxycoumarin (4-H) at B3LYP/6-31G* level of the theory.
2.638 Å and 2.696 Å, showed comparatively small deviations from the experimental values, 0.014 and 0.024 Å. The average deviations of the calculated bond lengths and bond angles at B3LYP/6-31G* are 0.81 % and 0.39 % respectively. The comparison of the calculated structural parameters of PhD and 4-H showed that in PhD differences were obtained in the O-H, =O and -O bond lengths. Due to participation in hydrogen bondings both hydroxyl and carbonyl groups in PhD showed longer O-H and =O bond lengths and shorter -O bond length: 0.971 Å 02/00 Å, 1.208 Å 1.229/1.234Å and 1.353 Å 1.327/1.330 Å. Using the same levels of theory, the geometry of D was fully optimized. As it is expected for the symmetrical HBs in D, the calculated O O distances do not differ as in the case of PhD. With the exception of some structural parameters in the exocyclic rings, the calculated parameters of D are very similar to that of PhD. 3.2. Estimation of the HB strengths in PhD and D. The strengths of the HBs formed, were estimated from the energies of different model structures. The global minimum structure, PhD1, was stabilized by one HB, O 10 -H 13 O 27, and the energy of this structure was used to estimate the second HB, O 28 -H 31 O 9. The O 28 -H 31 O 9 HB strength was evaluated from the PhD and PhD1energy difference. It was calculated of 52.32 kj/mol. By analogy, the energy of a model structure with only O 28 -H 31 O 9 HB, PhD2, was used to estimate the O 10 -H 13 O 27 HB in PhD; the calculated energy was 55.46 kj/mol. The HB energies thus obtained correlated with the calculated and experimental O O distances and predicted by 3.14 kj/mol stronger O 10 -H 13 O 27 HB. The total HB energy in PhD was estimated of 107.78 kj/mol. For D, a global minimum structure with both HBs ruptured, D1, was localized on the PES, and the total HB energy in D was calculated as a difference between the energy of D and the energy of D1, -101.79 kj/mol. Since the HBs in D are symmetrical, the single HB energy in D was calculated as a half of the total HB energy, -50.89 kj/mol. The value thus obtained predicted that the HB in D approached in strength the weaker HB in PhD (-52.32 kj/mol). The calculated HB energies in PhD (-55.46 kj/mol and 52.32 kj/mol) and in D (-50.89 kj/mol) pointed out to relative strong HB in the systems studied. The values
obtained showed greater HB stabilization than that obtained for normal HB (8-42 kj/mol) [3]. At the same time, they are below the values, obtained for some negative charge assisted O-H O hydrogen bonds, where the energy gain by HBs formation in the active site of enzymes could be up to 84 kj/mol, depending strongly on the nature (hydrophobicity, electrostatics, etc.) of the active site [3-5]. 3.3. Natural population analysis of PhD, D and 4-H. Natural population analysis data helped us to study the charge changes going from the parent compound, 4-H, to D and PhD. For PhD, less negative natural charges were found on the oxygen atoms, involved in HB in the upper exocyclic ring, O 10 and O 27, (-0.680, -0.637) as compared to the natural charges obtained on O 28 and O 9 in the lower exocyclic ring (-0.690, -0.654). The less negative charges on both the carbonyl and hydroxyl oxygen atoms in the upper exocyclic ring explained the shorter O 10 O 27 distance in comparison with the O 9 O 28 one. At the same time, the charge on H 13 is more positive (0.531) in comparison with that on H 31 (0.527). As it is expected, the hydroxyl oxygen atoms (O 10 and O 28 ) have more negative natural charges (-0.680 and -0. 690) as compared to the carboxyl oxygen atoms (O 9 and O 27 ) (-0.654 and -0.637). The results thus obtained suggested that a substantial electrostatic interaction is present in both HBs of PhD and from purely electrostatic arguments the O 10 -H 13 O 27 bonding in the upper ring is predicted to be stronger as compared to the O 28 -H 31 O 9 one. The carbonyl oxygen atoms in D have identical and larger charges (-0.651) as compared to 4-H (-0.570). The same results were obtained for the hydroxyl oxygen atoms (-0.671-0.687). Both hydroxyl and carbonyl oxygen charges in D have values, which are closer to those of the hydroxyl and carbonyl oxygen atoms involved in the weaker HB in PhD, O 28 and O 9 and these results confirmed the suggestion that the HBs in D approach the weaker HB in PhD. 3.4. Electron density (ρ b ) and its Laplacian at the bond critical points ( 2 ρ b ) of PhD, D and 4-H. The calculated electron density properties of PhD and D showed that both O H bondings have low ρ b, (ranging from 0.0443 to 0.0483), and positive 2 ρ b values (ranging from 366 to 524). These properties are typical for closed-shell interactions as HBs and indicate electrostatic character of the O-H O bondings [19]. As it was found from the calculated HB energies, the O-H O HBs in PhD are different in
strength and hence the calculated ρ b, and 2 ρ b, for the H-bonding in the upper exocyclic ring were higher (0.0483 and 524, respectively) than those in the lower exocyclic ring (0.0444 and 371). The results obtained are consistent with more electrostatic O-H O bonding and increasing bond strength in the upper exocyclic ring and they are in full agreement with the correlation between ρ b and 2 ρ b from one side, and the H-bonding strength from the other. This correlation was first reported by arroll and Bader [21] and confirmed later by other authors [22,23]. The HBs in D are symmetrical and the calculated ρ b, and 2 ρ b, for both O H bondings are obtained identical and comparable with the values of the weaker HB in PhD. The Laplacian at the critical points of the O-H bonds in PhD has different negative values, O 10 -H 13 = -1.557 and O 28 -H 31 = -1.576, thus pointing out that the O 10 -H 13 bond has lower covalent contribution (lower negative value). The lower covalent character of O 10 -H 13 bond as compared to that of O 28 -H 31 correlated with more electrostatic O 10 -H 13...O 27 HB. Identical and negative values for the Laplacian at both O- H bond critical points were found for D (-1.567) indicating no difference in the O- H O bondings in this compound. 3.5. Estimation of the HB strength in PhD and D from ν(o-h) vibrational shifts. The calculated vibrational shifts, ν (O-H), were obtained as a difference between the scaled wavenumbers of PhD or D and the corresponding scaled wavenumbers of 4- H. In the PhD IR spectrum, higher negative shift was obtained for the ν(o 10 -H 13 ), - 601 cm -1 (exp. shift, -602 cm -1 ) in comparison with the shift of ν(o 28 -H 31 ), -559 cm -1 (exp. shift, -560 cm -1 ). The higher red shift of ν(o 10 -H 13 ) is consistent with higher lengthening of O 10 -H 13 ( r(o 10 -H 13 ) = 0.031 Å) as compared to that of O 28 -H 31 ( r(o 28 - H 31 ) = 0.029 Å). The higher red shift of ν(o 10 -H 13 ) and the higher lengthening of O 10 -H 13 in PhD are in excellent agreement with the other molecular properties discussed in the previous sections (lower covalent character of O 10 -H 13 bond, more positive charge on H 13, smaller negative charges on O 10 and O 27, and shorter O 10 O 27 distance). All these properties together with the estimated HB energies ( 55.46 kj/mol and 52.32 kj/mol) pointed out that the HB in the upper exocyclic ring of PhD is stronger.
In D, both O-H lengthenings are comparable with the lengthening of the O 28 - H 31, involved in the weaker HB in PhD. In consequence, the negative shifts of ν(o 28 - H 31 ) and ν(o 10 -H 13 ) in D (-562 and -566 cm -1, respectively) were found comparable, with ν(o 28 -H 31 ) in PhD (-559 cm -1 ). The shifts obtained were in agreement with the calculated O O distances and HB energies and also predicted that both HBs in D are comparable with the weaker one in PhD. References [1] E.J. Valente and D.S. Eggleston, Acta ryst. 45 (1989) 785. [2] D.W. Hutchinson and J.A. Tomlinson, Tetrahedron 25 (1969) 2531. [3] B.Schiøtt, B.B. Iversen, G.K.H. Madsen and T.. Bruice, J. Am. hem. Soc. 120 (1998) 12117 (and references therein). [4] M.G.-Viloca, A.G..-Lafont and J.M. Lluch, J. Am. hem. Soc. 118 (1997) 1081. [5] A. Tomlinson, Ph.D. Thesis, University of Warwick (1968). [6] (a) P. Gilli, V. Bertolasi, V. Ferretti and G. Gilli, J. Am. hem. Soc. 116 (1994) 909; (b) G. Gilli F. Bellucci, V. Ferretti, V. Bertolasi, J. Am. hem. Soc. 111 (1989) 1023. J. [7] M.J. Frisch et al., Gaussian 98, A.&, Gaussian Inc., Pittsburgh, PA, 1998. [8] A.D. Becke, J. hem. Phys. 98 (1993) 5648. [9]. Lee, W. Yang, R.G. Parr, Phys. Rev. B37 (1998) 785. [10] P.R. Rablen, J.W.Lockman and W.L. Jorgensen, J. Phys. hem. A 102 (1998) 3782. [11] J.J. Novoa,.J. Sosa, J. Phys. hem. 99 (1995) 15837. [12] D.R. Mamann, Phys. Rev. B (1997) R10157. [13]. Maerker, P.v.R. Schleyer, K.R. Liedl, T.-K. Ma, M.Quack, M.A. Suhn, J. omp. hem. 18 (1997) 1695. [14] M. Lonzynski, D. Rusinska-Raszak, H.-G. Mack, J. Phys.hem. A 102 (1998) 2899. [15] K. Raghavachari, G.W. Trucks, J. A. Pople, M. Head-Gordon, hem. Phys. Lett. 57 (1989) 479. [16] R.F.W. Bader, S.G. Anderson, A.J. Duke, J. Am. hem. Soc. 101 (1979) 1389. [17] R.F.W. Bader, T.S. Slee, D. remer, E. Kraka, J. Am. hem. Soc. 105 (1983) 5061. [18] R.F.W. Bader, P.J. MacDougall, J. Am. hem. Soc. 107 (1985) 6788. [19] R.F.W. Bader, P.J. MacDougall,.D.H. Lau, J. Am. hem. Soc. 106 (1984) 1594. [20] (a) G. Bravic, J. Gaultier and. Hauw,. R. Acad. Sci. Ser., 267 (1968) 1790; (b) J. Gaultier and. Hauw, Acta ryst. 20 (1966) 646. [21] M.T. arroll, R.F.W. Bader, Mol. Phys. 63 (1998) 387. [22] J.A. Platts, S.T. Howard, B.R.F. Bracke, J. Am. hem. Soc. 118 (1996) 2726. [23] J. Luis Perez-Lustres, M. Mosquera, T. Klark, Phys. hem. hem. Phys. 3 (2001) 3569.
Hydrogen Bonded omplexes of Acetylacetone and Methanol: HF and DFT level Study Vassil B. Delchev Department of Physical hemistry, University of Plovdiv Paisij Hilendarski, Plovdiv 4000, Bulgaria Accepted for publication in Monatshefte für hemie Acetylacetone and its two tautomeric forms have been widely studied mainly because of the interest in the mechanism of keto-enol tautomerism in the gas phase (for isolated molecules) as well as in different solvents [1-4]. The solvent polarity influences the strength of the intramolecular hydrogen bond in the enol form and the composition of the acetylacetone tautomers. The extraordinary stability of the enol form is related to the formation of this bond and forms a cyclic pseudoaromatic structure [5,6]. Several experimental techniques and theoretical methods have been used for determination of the acetylacetone structure [7-9]. Up to now there is no clear idea about the possible interactions of acetylacetone with other organic molecules (e.g. methanol) by intermolecular hydrogen bonds. Therefore, complexes of the acetylacetone tautomers with methanol are reasonable models for this. Thus, the aim of this paper is: 1) to study comparatively at four theoretical levels the possible bonding between one acetylacetone molecule and one methanol molecule by intermolecular H-bonds; and 2) to clarify the details of the intermolecular interactions. This theoretical study was performed by means of ab initio calculations including density functional (DFT) and Hartree Fock (RHF) theory. RESULTS The optimized structures of all possible H-bonded complexes between acetylacetone tautomers and methanol are illustrated in Fig. 1. The calculated vibrational spectra of the complexes show that all =O vibrations are mixed with other motions of the molecule, some of them out of plane. For example, in the enol complexes the =O vibration is mixed with deformations of the methyl hydrogen atoms.
K6O E1H E5H E8O Fig. 1. Optimized structures of the H-bonded complexes For the system K6O one torsional vibration of the methyl hydrogen atoms was detected at 1798 cm -1 together with the characteristic =O band. According with Tayyari s investigations [10] the methyl stretching vibrations in the interval 3000 3850 cm -1 have very low IR-intensities. They should be active mainly in Raman
spectra of the systems. O-H stretching vibrations of the enol complexes restricted in the interval 2784 2913 cm -1 have quite high intensities. The same vibration in the pure acetylacetone enol form is at 2800 cm -1 in the gas phase and at 2875 cm -1 in the liquid phase [10]. In chloroform this vibration is at 2750 cm -1 as a broad asymmetric stretching band [11] and the carbonyl stretchings are at about 1550 cm -1 [11]. In spite of the fact that the experimental spectrum was measured in the liquid phase and all theoretical spectra refer to isolated systems we obtained very good correlations between experimental and theoretical frequencies with high correlation coefficients (given in brackets): 1) K6O: ν th = 0.947, ν exp + 23.48 (0.998); 2) E1H: ν th = 0.939, ν exp + 39.61 (0.996); 3) E5H: ν th = 0.910, ν exp + 83.29(0.993); and 4) E8O: ν th = 0.894, ν exp + 93.11 (0.991). The complex formation of the K6O and E1H systems is much more favorable than the complex formation of the E5H and E8H systems in the gas phase and in solvents. Further, these complexes (K6O and E1H) were found to have the lowest energies and the strongest (shortest) intermolecular hydrogen bonds. Moreover, the experimental IR spectrum gave bands similar to the calculated frequencies of the two complexes. The energy differences between the E1H complex and K6O, E5H, and E8O are 20, 6 and 16 kj mol -1. The first value gives the heat effect of the transformation K6O E1H (keto-enol tautomerism) in the gas phase. The keto-enol conversion between these two forms is enthalpically unfavored. It was found that the intermolecular H-bonds between the acetylacetone enol form and methanol causes a lengthening of the intramolecular H-bond, i.e. a decrease in the intramolecular H-bond energy. This influence is larger in the complexes whose intermolecular hydrogen bonds are from the O(acac)..H(meth) type (e.g. the E1H complex, in which this effect is the highest) and smaller in the E8O complex in which intermolecular H-bond is formed between H(8) and the methanol O-atom. References [1] Delchev VB, Mikosch H, Nikolov GSt (2001) Monatsh hem 132: 339 [2] Delchev VB (2000) Scientific works of Plovdiv University 29: 79 [3] Nikolov GSt, Markov P (1977/78) Annual of Sofia University 1: 109
[4] Nikolov GSt, Markov P (1981) J Photochem 16: 93 [5] Burdett JL, Rogers MT (1964) J Am hem Soc 86: 2105 [6] Rogers MT, Burdett JL (1965) an J hem 43: 1516 [7] amerman A, Mastropaolo D, amerman N (1983) J Am hem Soc 105: 1584 [8] Lowrey AH, George, D Antonio P, Karle J (1971) J Am hem Soc 93: 6399 [9] Pashkevich KI, Salutin VK, Postovski IY (1981) Russ hem Revs 50: 325 [10] Tayyari SF, Milani-nejad F (2000) Spectrochim Acta A 56: 2679 [11] Mavri J, Grdadolnik J (2001) J Phys hem A 105: 2045