Unrecognized Intramolecular and Intermolecular Attractive Interactions

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Unrecognized Intramolecular and Intermolecular Attractive Interactions between Fluorine-Containing Motifs and Ether, Carbonyl and Amino Moieties Kenneth B Wiberg, William F Bailey, and Kyle M. Lambert J. Org. Chem., Just Accepted Manuscript • DOI: 10.1021/acs.joc.9b00780 • Publication Date (Web): 01 Apr 2019 Downloaded from http://pubs.acs.org on April 8, 2019

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The Journal of Organic Chemistry

Unrecognized Intramolecular and Intermolecular Attractive Interactions between Fluorine-Containing Motifs and Ether, Carbonyl, and Amino Moieties Kenneth B. Wiberg,*, † William F. Bailey,*, ‡ and Kyle M. Lambert‡ † Department of Chemistry, Yale University, 275 Prospect Street, New Haven, CT 06520-8107, United States ‡ Department of Chemistry, University of Connecticut, 55 North Eagleville Rd, Storrs, CT, 06269-3060, United States

H3C F F O F H3C F polar non-polar • attraction energy > 7 kcal / mol! • significant charge transfer •unrecognized interaction

ABSTRACT: In order to elucidate to what extent Coulombic and other interactions contribute to the origins of contrasteric phenomena we have identified a significant, previously unrecognized interaction between fluorine-containing motifs and groups or molecules containing main-group heteroatoms. The axial conformers of both 2-methoxy- and 2-trifluoromethoxytetrahydropyrans preferentially adopt a rotameric arrangement in which the OCH3 and the OCF3 groups are gauche to the ring oxygen. Given that one would expect a repulsive Columbic interaction to exist between the electronegative fluorines of the CF3 group and the ring oxygen in this rotomeric orientation, this 1 ACS Paragon Plus Environment

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surprising result suggests that an attractive interaction exists between the CF3 group and the oxygen of the ring. The generality and origin of this interaction was examined using non-polar CF4 to probe intermolecular interactions with systems such as dimethyl ether, trimethylamine, trimethylphosphine and acetone. In each case there was an attractive interaction leading to formation of a complex. The attraction is not due to van der Waals forces. Rather, the fluorine lone pairs of the CF4 often act as an electron donor in these complexes leading to a transfer of charge between the reactants and formation of the complex. These previously unrecognized fluorineheteroatom interactions likely play a significant role in the context of understanding the binding interactions of medicinally relevant molecules or pharmaceuticals possessing fluorine-containing pharmacophores with their targets.

INTRODUCTION We recently reported the results of a coordinated experimental and computational reexamination of the anomeric effect which concluded that the phenomenon, while complex in origin, is, in the aggregate, mainly the result of attractive Coulombic interactions.1 While in the process of extending this study to an analysis of the rotameric arrangements of alkoxy groups at the anomeric carbon of glycosides (the so-called exo-anomeric effect),2 we had occasion to examine axial 2-methoxytetrahydropyran (2-OCH3 THP) at the MP2/6-311+G* level. As would be expected (Figure 1),2 the rotameric form in which the methyl group of the OCH3 moiety is gauche to the ring oxygen (1) was found to be 4.2 kcal/mol more stable than the anti form (2). We anticipated that this rotameric situation would be reversed in the case of a THP bearing an axial 2OCF3 group due to Coulombic repulsion between the fluorines of the CF3 group and the proximate ring oxygen in the gauche rotamer (3). This proved not to be the case. To our surprise, geometry

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optimization of the anti rotamer of the axial 2-OCF3 compound (4) proceeded to give the syn isomer (3): the anti form (4) was found to have one imaginary frequency. Clearly, there is something unexpected in the nature of the apparently attractive interaction between the CF3 group and the ring oxygen in 3.

H

O CH3

O

H

O

O

O

O

H

O

1

O

H

O CF3

CH3

CF3 O

3

CH3

H3C

H O

H

O O O

2

CF3 H

F3C

H O

O

4

Figure 1. Rotameric conformations of axial 2-OCH3 THP and axial 2-OCF3 THP.

RESULTS AND DISCUSSION In order to better understand the nature of the CF3 – ring oxygen interaction, rotational profiles of the axial conformers of 2-OCH3 THP and of 2-OCF3 THP were calculated at the MP2/6-311+G* level. In order complete this analysis, the corresponding rotational profiles of equatorial conformers were also examined. The torsional angle was taken as H(2)–C(2)–O–C and the results are shown in Figure 2. When examining these profiles, it should be noted that the rotamers labeled syn correspond to conformations in which the CH3 or CF3 group is gauche to the ring oxygen, the anti label denotes the conformation in which these groups are gauche to the ring carbon, and the minima labeled trans correspond to the conformation in which the groups are anti with respect to H(2). 3 ACS Paragon Plus Environment


Figure 2. Rotational profiles of 2-OCH3 THP and 2-OCF3 THP.

The plots for the corresponding 2-OCF3 THP isomers are similar to those of the OCH3 compounds, but there are significant differences. Whereas the anti rotamers are true minima with the CH3O substituent, these rotamers are merely small bumps on the surface for the OCF3 isomer. Consequently, it becomes clear why attempted calculations for these anti rotamer forms led 4 ACS Paragon Plus Environment

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directly to the syn isomers. The most significant observation is that the syn conformations are the global minima for both the OCH3 and OCF3 substituted tetrahydropyrans. An attempt was made to gain a clearer understanding of these rotational profiles by examination, using the Hirshfeld charges,3,4 of charge distributions within the molecules as a function of rotation. However, in addition to changes in the charge distributions in the 2-alkoxy substituents, there were, as might be expected, also changes for the ring atoms and it was not clear which charge differences were of importance. We should note at this juncture that a hyperconjugative interaction,5 involving electron delocalization of a 2p-type lone pair on the oxygen of a 2-alkoxy substituent (often, as here (Figure 1)) rendered as an sp3 hybridized lone pair) into the adjacent antibonding σ*-orbital of the ring C(2)–O unit, may well contribute to the relative stability of the syn rotamers in both axial and equatorial isomers of 2-OCH3 or 2-OCF3 THP. However, it is not at all clear how that this factor alone can account for the relative instability of the anti rotamer of a 2-OCF3 group or the broad global minima found in the rotational profile for this group. In an attempt to investigate the relevant interactions responsible for the behavior of the THP compounds,

we

examined,

also

at

the

MP2/6-311+G*

level,

a

simpler

model:

methoxytrifluoromethoxymethane (CH3OCH2OCF3), (5). The calculated the rotational profile for the O–C–O–CF3 torsion angle from 180 to 0°, with the CH3O–C–O–C torsion angle set to a normal value of 60°, is shown in Figure 3. The minimum energy conformation at 70° is ~1.5 kcal/mol lower in energy than the 180° conformer. Moreover, as illustrated in Figure 4, this lowest energy rotamer of 5 is structurally quite similar to the minimum energy syn conformation of axial 2-OCF3 THP (3). In short, the model structure (5) includes the essential structural features of 3.



10

8

6

4

2

0

-2 0

36

72

108

144

180

torsion angle Figure 3. Rotational profile for methoxytrifluoromethyoxymethane.


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Figure 4. Minimum energy structures of compounds 5 (top) and 3 (bottom) showing the alignment of a C–F bond with a C–O bond.

A detailed analysis of variations in charge distribution within a molecule as it undergoes a conformational change is often able to provide significant additional information.1 We have recently demonstrated that the Hirshfeld approach for determining atomic charges,3 particularly for hydrogens, is superior to other methods because the results of a Hirshfeld analysis are uniquely able to reproduce experimental data.4 The Hirshfeld charges for the 70° minimum energy conformation of 5 (Figure 4) and the corresponding 180° conformer of 5 are compared in Table 1.



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Table 1. Comparison of Hirshfeld Charges, q (e), for the 180° and 70° Rotamers of 5.

a

Atom or group

180°

70°

changea

CH3

0.1263

0.1203

–0.0060

O(4)

–0.1580 –0.1652 –0.0072

C(1)

0.1073

0.1063

–0.0010

H(2)

0.0365

0.0472

0.0107

H(3)

0.0518

0.0492

–0.0026

O(5)

–0.1632 –0.1675 –0.0043

CF3

–0.0003

0.0096

0.0099

Charge in the 70° conformer – charge in the 180° conformer.

An examination of the data in Table 1 demonstrates that there are significant changes in charge distribution within molecule 5 on going from a torsional angle of 180° to the minimum energy structure at 70°. The two entries in Table 1 that lose significant charge on going to the 70° conformer are the CF3 group, that donates ~0.010e, and H(2) that loses ~0.011e. The substantial loss of charge at H(2) would result in an attractive, stabilizing Coulombic attraction for the neighboring oxygens, O(4) and O(5), that are seen to gain 0.007e and 0.004e on going from a torsional angle of 180° to the minimum energy 70° conformer. The loss of charge from the CF3 group was an unexpected result of this analysis. Indeed, it may seem surprising that a CF3 group could function as an electron donor in this context. However, the lone pairs on fluorine are relatively compact, in comparison to those on a nitrogen or an oxygen,


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and the two electrons in the lone pair would have a large repulsion energy that may be reduced by transfer of some electron density from fluorine to an empty orbital of a nearby atom. Etiology aside, the results suggest that there is a potentially important, unrecognized interaction of some sort between CF3 groups and oxygens in the intramolecular situations discussed above. In order to see if this is a general occurrence, the analysis was extended to investigate intermolecular complexes between neutral heteroatom-containing molecules and carbon tetrafluoride (CF4). As the results discussed below demonstrate, the interaction of CF3 groups and heteroatoms appears to be a rather general phenomenon. The presence of fluorine atoms in organic molecules has long been known to profoundly affect the physical, chemical and biological properties of such materials.6,7 Moreover, anionic complexes between fluoride or chloride and CF4 have recently been reported.8 An important source of stabilization of these anionic intermolecular complexes is the reduction of the electrostatic energy of F– or Cl– as they transfer some charge to the CF4 molecule. In order to gain some insight concerning the apparently attractive interaction between an oxygen and a CF3 group, we have carried out calculations at the HF and MP2 levels using both a medium-sized basis set (6-311+G*) and a large basis set (aug-cc-pVTZ) for complexes between carbon tetrafluoride, a surrogate for nonpolar, fluorine-containing moieties, and the following representative heteroatom-containing compounds: dimethyl ether, acetone, trimethylamine and trimethylphosphine. We reasoned that the results of the HF calculations would indicate if there is an intrinsic stabilization upon forming a complex; the MP2 results would reveal the effect of electron correlation. The smaller basis set may lead to a basis set superposition error, but this should be essentially eliminated using the larger basis set. The results of these calculations,



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tabulated as the difference in energy in kcal/mol between the complexes and the isolated reactants, are given in Table 2.

Table 2. Stabilization Energies in kcal/mol for the CF4 Complexes. complex

HF (6-311+G*)

HF (aug-cc-pVTZ)

MP2 (6-311+G*)

MP2 (aug-cc-pVTZ)

(CH3)2O – CF4

1.28

0.60

3.18

2.79

(CH3)2C=O – CF4

1.14

0.66

2.64

2.11

(CH3)3N – CF4

1.33

0.61

3.80

3.30

(CH3)3P – CF4

0.61

0.25

2.18

1.46

The results demonstrate that there is little stabilization of the complexes at the HF level and that the smaller basis set leads, as expected, to a basis set superposition error. The MP2 calculations gave considerably larger stabilization energies, and here again the smaller basis set leads to a basis set superposition error. The main difference between the computed HF and MP2 structures is the O – CF4 and N – CF4 distances in the complexes. For example, in the case of the (CH3)2O complex, displayed in Figure 5, the O – CF4 distance is reduced from 3.552 Å in the HF structure to 3.14 Å in the MP2 structure. This is as expected since electron correlation reduces the repulsion between the two partners. The following discussion refers to the MP2/aug-cc-pVTZ results. The structure of the (CH3)2O – CF4 complex, which has Cs symmetry, is displayed in Figure 5. The angle formed from the midpoint between the carbons of (CH3)2O, the oxygen and the carbon of CF4 is ~124°. This is the angle typically associated with the oxygen lone pairs.9 In the complex, two of the three C–F bonds of CF4 closest to the oxygen of the ether are shorter, and the remote


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C–F bond is longer, than are the C–F bonds in the non-complexed CF4 molecule. The structural changes are summarized in detail in the Supporting Information.

Figure 5. (CH3)2O – CF4 complex: the geometrical arrangement of the ether with respect to CF4 is shown on the left; the Cs symmetry of the complex is illustrated on the right.

The formation of a complex between (CH3)2O and CF4 must involve a shift in charge density. The simplest proposal would posit that the oxygen of the ether donates charge to the CF4 group and that much of that charge goes to the remote C–F bond. However, as revealed by examination of atomic charges in the reactants and the complex, it is not quite that simple. The Hirshfeld charges for (CH3)2O, CF4 and the complex, are listed in Table 3.



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Table 3. Hirshfeld Charges for Me2O and CF4 and the Complex between them, q (e). reactant

O of Me2O

–0.1733

–0.1716

0.0017

each Me of Me2O

0.0866

0.0773

–0.0093

C of CF4

0.3727

0.3736

0.0009

each of 3 proximal F

–0.0930

–0.0852

0.0078

distal F

–0.0930

–0.1010

–0.0080

a

complex

changea

atom or group

Charge in complex – charge in isolated reactant.

The data summarized in Table 3 demonstrate that the sum of the (CH3)2O charges in the complex [ q = –0.1716 + 2(0.0773)] is –0.0171e and that for CF4 in the complex [ q = 0.3736 + 3(–0.0852) + (–0.1010)] is + 0.0171e. Thus, there is indeed a net shift of charge density upon formation of the complex. However, contrary to our initial expectations, the shift in charge involves a net transfer of charge density from CF4 to (CH3)2O. One might reasonably ask: how can that be? The overall charge in the methyl groups of (CH3)2O in the complex is less positive than that of these groups in the isolated ether. Evidently, upon formation of the complex, charge density is transferred from the three proximal fluorines (Figure 5) to the methyl groups. This is likely a consequence of a transfer of density from the non-bonded electrons of the fluorine atoms to the antibonding *-orbital of the C–O unit that has the larger coefficient at the carbon. With the Hirshfeld charges in hand, it is possible to estimate the Coulombic repulsion between the negatively charged oxygen of the ether and the three nearby, negatively charged, fluorines of CF4 in the complex.10 This rather rudimentary calculation suggests that the Coulombic repulsion between the ether and CF4 is approximately 4.9 kcal/mol. Summation of the calculated stabilization 12 ACS Paragon Plus Environment

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of the complex from Table 2 (2.8 kcal/mol) and the repulsion that must be overcome in forming this complex (~5 kcal/mol) gives ~7.9 kcal/mol as the total attraction energy for the complex. Essentially the same picture emerges from an analysis of the complex between (CH3)3N and CF4. The computed structure of this complex, which has C3v symmetry, is depicted in Figure 6 and the Hirshfeld charges in the reactants and in the complex are summarized in Table 4. The nitrogen of the amine, which is more basic than the oxygen of (CH3)2O, transfers more electron density to CF4 than does the ether. However, the most important effect of complex formation is a transfer of 0.003e from the fluorines of CF4 to each of the methyl groups of the amine upon formation of the complex. Overall, there is a net transfer of 0.0036e from CF4 to the amine.

Figure 6. (CH3)3N – CF4 complex.

Table 4. Hirshfeld Charges for (CH3)3N and CF4 and the Complex between them, q (e). 13 ACS Paragon Plus Environment


Atom or group

reactant

N of (CH3)3N

complex

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changea

–0.1042

–0.0989

0.0053

each Me of (CH3)3N

0.0347

0.0317

–0.0030

C of CF4

0.3721

0.3716

–0.0005

each of 3 proximal F

–0.0930

–0.0878

0.0052

distal F

–0.0930

–0.1046

–0.0116

aCharge

in complex – charge in isolated reactant.

The structure of the acetone – CF4 complex is depicted in Figure 7. The angle described by the carbonyl carbon to carbonyl oxygen to the CF4 carbon is 132° indicating that the oxygen lone pair is, as expected, in the direction of the CF4 carbon. The structure differs in another significant way from those of the ether and amine complexes discussed above: one of the methyl groups of the acetone molecule in the acetone complex is unavoidably much farther from the CF4 than is the other. As a result, it might be anticipated that any charge transfer from CF4 upon formation of the complex would be appreciable less than in the previous cases.


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Figure 7. Acetone – CF4 complex.

The Hirshfeld charges for the reactants and the acetone – CF4 complex are summarized in Table 5. All four fluorines of CF4 in the complex are distinct; for this reason, the atoms in the complex have been numbered, as shown in Figure 5, to permit discussion of the charges on the four nonequivalent fluorines. Inspection of the charges tabulated in Table 5 demonstrates that there is some charge transfer from the carbonyl oxygen to the carbon of CF4 and there is significant charge transfer from F(12) and F(13) to the proximal methyl carbon [C(3)] but not to the more remote methyl carbon [C(7)]. However, overall, there is a net transfer of 0.0055e charge from acetone to CF4 upon formation of the complex. This is the reverse of the charge transfer found upon formation of a complex between CF4 and the ether or the amine.



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Table 5. Hirshfeld Charges for Acetone and CF4 and the Complex between them, q(e). atom or group

aCharge

reactant

complex

changea

O of (CH3)2C=O

–0.2443

–0.2368

0.0075

C of (CH3)2C=O

0.1649

0.1666

0.0017

proximal C(3)b of (CH3)2C=O

0.0397

0.0341

–0.0056

distal Me(7)b of (CH3)2C=O

0.0397

0.0416

0.0019

C of CF4

0.3721

0.3705

–0.0016

F(12)b

–0.0930

–0.0894

0.0036

F(13)b

–0.0930

–0.0900

0.0030

F (14)b

–0.0930

0.1046

–0.0114

F(15)b

–0.0930

–0.0920

0.0010

in complex – charge in isolated reactant. b Atoms are numbered as depicted in Figure 7.

It was recognized at the outset of this investigation that the intermolecular complexes CF4 with neutral, heteroatom-containing molecules discussed above may simply be a consequence of van der Waals forces. To explore this possibility, we examined the complex between CF4 and trimethylphosphine. The structure of this complex is depicted in Figure 8. If van der Waals forces are the principal factor responsible for complex formation, the complex between CF4 and (CH3)3P would be more stable than the corresponding amine complex because phosphorus is more polarizable than nitrogen. Alternatively, if the interaction of the CF4 molecule with a lone pair on the heteroatomic partner were of central importance to the formation of a complex, the stabilization of the complex between (CH3)3P and CF4 would be weaker than that of the amine because phosphorus is less basic than nitrogen. It was found (Table 2) that the (CH3)3P – CF4 complex has


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a stabilization energy of only 1.46 kcal/mol while that of the (CH3)3N – CF4 complex is 3.30 kcal/mol. Clearly, van der Waals attraction is not the major factor leading to the stability of the complexes described above. Indeed, the (CH3)3P – CF4 complex (Figure 8), is structurally quite different than that of the (CH3)3N – CF4 complex: the phosphine complex lacks the C3v symmetry of the amine complex (Figure 6) and the distance between the phosphorus and the CF4 carbon (~4 Å) is considerably longer than is the corresponding distance in the amine complex (~3.2 Å).

Figure 8. (CH3)3P – CF4 complex.

It might be noted at this juncture that the complex formed from fluoride ion and CF48 is more stable than any of the intermolecular complexes described above. An analysis of the Hirshfeld charges of this anionic complex suggest that approximately 15% of the fluoride charge is donated to the CF4 upon formation of the complex and this charge redistribution leads to significant reduction in the electrostatic energy. Clearly, the formation of intermolecular complexes of the sort discussed above results in loss of translational entropy. Consequently, the free energy change upon formation of intermolecular



complexes will be unfavorable. Of course, intramolecular interactions, as in the trifluoromethoxy substituted 2-OCF3 THP molecule (3), will not suffer from this unfavorable this entropy effect. Nevertheless, the finding of unanticipated stability in complexes between CF4 and molecules containing oxygen and nitrogen suggests, inter alia, that stabilizing attractions of this sort may well be an important, albeit unrecognized, factor at play in the interaction of fluorine-containing pharmaceuticals with receptor sites.11

CONCLUSIONS The results described above provide the first evidence, to our knowledge, for the presence of significant attractive interactions between groups or molecules containing main-group heteroatoms such as oxygen or nitrogen and CF4 or CF3 – containing groups. The attraction is apparently not a simple manifestation of van der Walls forces; rather, the interaction appears to be principally the result of a transfer of charge between the reactant lone pair and the positively charged carbon, and back-bonding from the fluorine lone pairs into the C-O* bond of the methyl groups. It remains to be determined whether intramolecular interactions of this sort are at play in more structurally complex molecules. Calculations: All calculations were carried out using Gaussian-16.12 The structures presented in the Figures were rendered using CYLview.13

ASSOCIATED CONTENT Supporting Information The Supporting Information is available free of charge on the ACS Publications website


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Tables of atomic coordinates and energies, of the complex energies, and of the Hirshfeld charges for reactants and complexes AUTHOR INFORMATION Corresponding Authors †E-mail; [email protected] ‡E-mail: [email protected] ORCID Kenneth B. Wiberg: 0000-0001-8588-9854 William F. Bailey 0000-0001-9159-0218 Kyle M. Lambert: 0000-0002-8230-2840 Present address (KBW) 85 Central Ave., Apt A404, Needham, MA 02492. Notes The authors declare no competing financial interest

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(9)

Legon, A. C.; Millen, D. J. Angular Geometries and Other Properties of Hydrogen-bonded Dimers: A Simple Electrostatic Interpretation of the Success of the Electron-pair Model. Chem. Soc. Rev. 1987, 16, 467–498.

(10) The Columbic interaction is given by E = (qaqb)/rab where the q’s are the atomic charges in e, and rab is the distance given in Bohrs. The result may be multiplied by 627.51 to give the energy in kcal/mol. (11) Ilardi, E. A.; Vitaku, E.; Njardarson, J. T. Data-Mining for Sulfur and Fluorine: An Evaluation of Pharmaceuticals To Reveal Opportunities for Drug Design and Discovery. J. Med. Chem. 2014, 57, 2832–2207. (12) Gaussian 16 Revision A.03. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Peterson, G. A.; Nakatsuji, H.; Li, X.; Caricto. M.; Marenich, A. V.; Bloino, J.; Janesko, B. G.; Gomperts, R; Mennucci, B.; Hratchian, H. F.;.Ortiz, J. V.; Izmayiov, A. F.; Sonnenberg, J. L.; Williams-Young, D.; Ding, F.; Lipparini, F.; Egidi, F.; Goings, I.; Peng, B.; Petrone, A.; Henderson, T; Ranasinghe, D.; Zakrzewski, V. G.; Gao, J.; Rega, N.; Zheng, G.; Liang, W.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Throssell, E.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.J.; Heyd, J. J.; Brothers, E. N.; Kudin, K. N.; Staroverov, V. N.; Keith, T. A.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A. P.; Burant, J. C.; Iyengar, S.S.; Tomasi, J.; Cossi, M.; Millam, J.M.; Klene, M.; Adamo, C.; Cammi, R.; Ochterski, J. W.; Martin, R. L.;



Morokuma, K.; Farkas, O.; Foresman, J. B.; Fox, D..J. Gaussian, Inc., Wallingford, CT, 2016. (13) Legault, C. Y. CYLview 1.0b; Université de Sherbrook, Sherbrook, QC, Canada, 2009; http://www.cylview.org.


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Unrecognized Intramolecular and Intermolecular Attractive Interactions

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