PMC:4264129 / 1419-21407 JSONTXT

{"target":"https://pubannotation.org/docs/sourcedb/PMC/sourceid/4264129","sourcedb":"PMC","sourceid":"4264129","source_url":"http://www.ncbi.nlm.nih.gov/pmc/4264129","text":"1. Introduction\nConformational preferences can cause non-contiguous atoms within an isolated molecule to become similarly close neighbors. These spatial arrangements may be driven by favorable electrostatic interactions or by the special case where three of such atoms form a so-called “hydrogen bond” (H-bond). Although the situation becomes more complicated when the molecular structure is considered within a solution environment, these same two factors remain important to also drive additional intermolecular interactions now possible between solute molecules themselves and with the solvent molecules as partners. Focusing herein on hydrogen bonding, it can be noted that, despite a decades-long endeavor to define the H-bond, this key arrangement still cannot be considered to be resolved with full consensus. The 2011 IUPAC recommendations provide a definition [1] that can be used as the basis for critical evaluation of a number of structural that will be examined in Section 3 of this review. The recommendations state:\n“The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation.\nA typical hydrogen bond may be depicted as X–H…Y–Z, where the three dots denote the bond. X–H represents the hydrogen bond donor. The acceptor may be an atom or an anion Y, or a fragment or a molecule Y–Z, where Y is bonded to Z. In some cases, X and Y are the same. In more specific cases, X and Y are the same and X–H and Y–H distances are the same as well leading to symmetric hydrogen bonds. In any event, the acceptor is an electron rich region such as, but not limited to, a lone pair of Y or π-bonded pair of Y–Z. The evidence for hydrogen bond formation may be experimental or theoretical, or ideally, a combination of both. Some criteria useful as evidence and some typical characteristics for hydrogen bonding, not necessarily exclusive. The greater the number of criteria satisfied, the more reliable is the characterization as a hydrogen bond.” A footnote (F1) emphasizes that “…there will be borderline cases for which the interpretation of the evidence might be subjective. In any case, there should be no gross deviations from the above-mentioned criteria.”\nSix criteria and several characteristics are also included in the article, and eight points, as footnotes, help the reader to interpret the used terms. The first criterion for a hydrogen bond claims: “The forces involved in the formation of a hydrogen bond include those of an electrostatic origin, those arising from charge transfer between the donor and acceptor leading to partial covalent bond formation between H and Y, and those originating from dispersion.” It reveals from the specification in footnote F2 that “Attractive interactions arise from electrostatic forces between permanent multipoles, inductive forces between permanent and induced multipoles, and London dispersion forces. If an interaction is primarily due to dispersion forces, then it would not be characterized as a hydrogen bond.”\nAccording to the classification of Grabowski [2], the stabilization energy from weak to moderate hydrogen bonds is 4–63 kJ/mol. The hydrogen bonds in the present survey generally reside in “weak” to at most “moderate” categories. For the molecules analyzed in section 3, the X and Y atoms are separated by at least two atoms along the covalently bonded substructure path. A basic requirement for the formation of a hydrogen bond is a short distance between the H and Y atoms. If the X and Y atoms are O or N, the X–H covalent bond is generally polarized, and the Y atom carries an electron lone-pair pointing toward the H atom. Through formation of a H-bond between the indicated atoms, dispersion forces are much less important than the electrostatic interactions and the charge transfers.\nIn a special case, Y symbolizes an aromatic ring with its electron cloud favorably interacting with a positively polarized H atom. This sort of hydrogen bond is called an H…π interaction. For X–H…Y with X = C or with X, Y = S or P, as well as for the H…π interaction, the role of the dispersion forces increases in comparison to the cases where the H-bond formation is principally related to electrostatic and charge-transfer effects. A situation where the dispersion interactions are apparently important will be discussed as part of the conformational analysis of tyramine in Subsection 3.3.1\nThe third criterion (E3) on the list of the IUPAC recommendations says: “The X–H…Y angle is usually linear (180°) and the closer the angle is to 180°, the stronger is the hydrogen bond and the shorter is the H…Y distance.” Two important footnotes were added to this criterion. “The X–H…Y hydrogen bond angle tends toward 180° and should preferably be above 110° (F4).” “Historically, the X to Y distance was found to be less than the sum of the van der Waals radii of X and Y, and this shortening of the distance was taken as an infallible indicator of hydrogen bonding. However, this empirical observation is true only for strong hydrogen bonds. This criterion is not recommended. In most cases, the distance between H and Y are found to be less than the sum of their van der Waals radii. It should be noted that the experimental distances are vibrational averages and would differ from such distances calculated from potential energy minimization. (F5)”.\nThus, as revealed by the quoted text, no H…Y or X…Y distance has been strictly defined for the distance of a H-bond, nor has a strict lower limit for the X–H…Y angle has. On the other hand, the X…Y distances for the different intramolecular H-bonds could represent borderline cases with values equal or slightly larger than the sums of the van der Waals radii. Likewise, in cases when a H-bond can form a five-member ring arrangement (Figure 1), the X–H…Y bond angles could be close to or even less than 110°.\nFigure 1 The figure shows the projection of the heavy-atom skeleton onto the X–H…Y plane for cases where H-bonding can result in a: (a) Five-member ring; (b) Six-member ring; or (c) Seven-member ring. Another point deserves consideration when qualifying whether a bond meets the criteria of being a H-bond. The atoms-in-molecules (AIM) theory of Bader and Popelier [3,4] identifies the H-bond in a topological manner. The theory can be applied to find bond critical points (BCP) and to analyze them in terms of electron densities and their Laplacian. Qualification is primarily related to the existence of a bond critical point of (3, −1) type, but there are seven more features to be met [5]. One of them is the mutual penetration of the hydrogen and the acceptor atoms. This characteristic of a H-bond guarantees that polar X–H and Y groups cannot form an intramolecular hydrogen bond if they are far from each other in the space. For a rigid system, e.g., 1,4-dihydroxy benzene, the chemical structure itself prevents the penetration. For saturated chain systems with conformational flexibility, an extended form like the trans conformation for the X–CH2–CH2-Y moiety prevents the necessary closeness of the polar groups (Figure 2). Consequently, no (3, −1) BCP could be found for the above structures.\nThe IUPAC recommendations, however, do not include the (3, −1) BCP as a necessary criterion for the existence of a H-bond. Instead, this feature is considered (C6) only a characteristic of a H-bond on the basis of theoretical results: “Analysis of the electron density topology of hydrogen-bonded systems usually shows a bond path connecting H and Y and a (3, −1) bond critical point between H and Y.”\nThis C6 point in combination with F1 above is very important in understanding several computational results. Klein [6] did not find a (3, −1) BCP regarding a O–H…O intramolecular hydrogen bond for 1,2 diols in their optimized geometries. Mandado et al., [7] found the (3, −1) bond critical point missing for gas-phase 1,2-dihydroxybenzene (catechol) with the geometry optimized at the B3LYP/6-311++G** level. The existence of this BCP would have indicated an O–H…O intramolecular hydrogen bond by the AIM theory. A slight distortion of the optimized catechol geometry, however, led to the appearance of the (3, −1) BCP. Thus, this molecule may present a borderline case for a H-bond (see F1 above).\nThis characterization of an intramolecular H-bond is largely retained for the case of an intermolecular H-bond with the basic difference that the X–H covalent bond and the Y atom (aromatic ring) are elements of two different molecules. In this case, the two species have to approach each other appropriately in space. Thus, whereas the intramolecular H-bond is a feature of a single molecule in a favorable conformation, the intermolecular H-bond between two molecules emerges only within a specific H…Y distance range. Accordingly, H-bond qualification at a separation corresponding to the sum of the van der Waals radii again becomes problematic. Regarding the X–H…Y bond angle, the values for an intramolecular and intermolecular bond could differ considerably. For the latter, calculations predict a slightly bent H-bond of about 160°–180° in the gas-phase unless there is an additional geometric constraint. The computational result is reasonable: the geometry optimization seeks a structure with minimized strain between the two species. On this basis, the favorable X–H…Y arrangement is close to linear. This conclusion refers only for isolated pairs, mostly existing in the gas phase. Alternatively, the crystalline phase environment can strongly affect the H-bond geometry [8].\nThus some points above suggest that borderline cases are conceivable, whereas no distance limit was provided in the IUPAC definition. On the other hand, consideration of the sum of the van der Waals radii, as an upper limit is not recommended. Indeed, what are the relevant van der Waals radii? Bondi presented a table for mean values [9], but are these values always relevant in any molecular environment? Could there be a H-bond if the sum of the considered van der Waals radii is almost equal to the H…Y distance?\nThese problems (and probably a number of others) underscore that the definition for a H-bond is not a closed chapter within the field of theoretical chemical research. Recently, Weinhold and Klein [10] published a paper with a detailed list of the former theoretical activities that have been directed toward this topic. The authors proposed a new H-bond definition in partial agreement with the present IUPAC recommendations. It is noteworthy that a topological requirement was also not put forward.\nIn a recent paper by Contreras-García et al. [11], the authors note that the density values at the H-bond critical point cannot be used to identify the most stable geometry of a complex. This statement is in accord with the results from calculations performed by Klein and Mandado et al. [6,7], although the latter also found that a (3, −1) BCP could be identified for 1,2-dihydroxy benzene upon a small geometry distortion, which suggests that the optimized and intramolecularly H-bonded structures are not too different. For a more effective analysis of non-covalent interactions, Contreras-García et al. [11] developed a new index (non-covalent interactions, NCI) by utilizing the reduced density gradient. Although the method is related to the AIM approach, the NCI features are tied to the critical points of the density gradient field. Use of the reduced density gradient facilitates the account for local density inhomogenities. The NCI isosurfaces reveal the connections of the critical points in the real space, which can form superbasins. Accordingly, as the authors claim, “ring or cage points are sometimes a better reference for understanding bond strength than bond point themselves.”\nThen the question remained in this review is: how to identify a hydrogen bond? The problem is more sensitive regarding the formation of an intramolecular H-bond because computations indicate that the (3, −1) BCP can generally be found for intermolecular hydrogen bonds. An important example is the microsolvation of a reference molecule by a few protic solvent molecules when the latter form a bridge within the intramolecular X–H…Y region [6,12,13,14,15,16]. Water and methanol can arrange in a manner that even one solvent molecule creates two intermolecular hydrogen bonds in the standard form of O–H…Y and X–H…O. In both arrangements, the H-bond distances can be of favorable length and the bond angles for the newly formed hydrogen bonds can be much closer to 180° than that for the unsolvated reference “solute”. Corresponding (3, −1) BCP’s have been found for the 1,2-ethanediol monohydrate [6] and for the 2,2,2-trifluoroethanol:water 1:1 complexes [16].\nIntramolecular H-bonds in the gas phase will be accepted in this review on the basis of experimental studies. For distance considerations, the values provided by Grabowski [2] will be utilized. The X…Y donor-acceptor separations for strong H-bonds with energy of 63–167 kJ/mol were accepted by Grabowski as 240–255 pm for O–H…O systems, 250–260 pm for the N–H…O bonds and 260–270 pm for the N–H…N interactions. No range was provided for the important O–H…N bonds, for which a characteristic distance of about 260 pm has been estimated here. The H-bonds in the present survey generally fall into the category “moderate”. Accordingly, the X…Y distances can be assumed to be at somewhat larger values than provided by Grabowski.\nAn important experimental feature of a H-bond is the shift of the X–H stretching frequency. In the case of “proper” hydrogen bonds, the frequency decreases and is called a red shift. Most of the H-bonds belong to this category [17]. In some cases, however, the X–H frequency increases upon H-bond formation. This phenomenon is called a “blue” or “improper shift”. The quantum mechanical comparison and the related explanation were summarized by Hobza and Havlas [18].\nThe shift in stretching frequency is related to the increase and the decrease of the X–H bond length in cases of the red- and blue-shifts, respectively. The change of X–H bond length is related to a charge transfer from the acceptor to the donor molecule in the H-bonded complex, which can be ascertained by means of NBO analyses [19]. In the case of the red-shift, some charge is transferred from the lone pair of the Y atom to the antibonding X–H orbital of the donor molecule. As a result, the X–H bond length increases and its stretching frequency decreases. Alternatively, blue-shift was noticed, for the Cl−…H3CBr system, as well as others. In this complex, the charge is transferred from the anionic acceptor to the antibonding orbital of the C–Br bond. The C–Br bond elongates followed by a geometry reorganization of the H3CBr molecule. In its new geometry, the C–H bond becomes shorter providing the basis for the frequency increase and hence the blue-shift in the spectrum. An important point of the review by Hobza and Havlas is the discussion of the (3, −1) BCP’s, which were found both for the proper and improper intermolecular hydrogen bonds. Other AIM criteria for a H-bond were also met for the systems exhibiting blue-shift of the vibrational frequency.\nFigure 2 Structures with an intramolecular hydrogen bond for: (1) 1,2-Ethanediol; (2) Salicylic acid; (3) Hydroxy-benzoic acid; and (5) β-Alanine zwitterion. Conformations 2, 4, 6 prevent the formation of the intramolecular H-bond and are open for forming intermolecular hydrogen bonds. For many systems studied below, AIM analyses were not found during literature searching. Furthermore, even when such calculations are performed, there remains the possibility of not finding a (3, −1) BCP, as happened to the optimized geometry of 1,2-ethanediol [6,7]. Thus the present author does not signify a H-bond to be present or not present based upon the existence of a BCP. This stance is supported by the allowable borderline systems in the IUPAC definition and by the conclusions from the NCI analyses [11] regarding energy-minimized structures. The existence of an intramolecular H-bond will be accepted if the experimentally derived H…Y distance is smaller than the sum of the van der Waals radii and/or a meaningful shift in the X–H stretching frequency was experimentally recorded.\nFor a number of isolated molecules, experiments predict (X) H…Y separations within the 200–250 pm range, with van der Waals radii of 120, 155, and 152 pm for H, N, and O, respectively [9]. For five-member saturated rings (Figure 1), the conformation corresponds to a X–C–C–Y gauche arrangement. Even if a (3, −1) BCP is missing, it is conspicuous that this conformation is the most stable one for many molecules in the gas phase. The aim of this review then becomes to consider the solvent effect on the maintenance or modification of the experimentally found gas-phase conformation while leaving the possibility open for forming an intramolecular H-bond. A gauche to trans transformation of the X–CH2–CH2–Y moiety would definitely disrupt an intramolecular H-bond (Figure 2). The intramolecular H-bond also is disrupted upon rotation of 180° about the C–O bond for species 3.\nEven if the gauche structure for the XCCY moiety is maintained, the intramolecular H-bond associated with a H–X–C–C gauche arrangement would be disrupted upon rotation about the X–C bond resulting in a H–X–C–C trans conformation (Figure 3). In the case of a six-member intramolecular H-bond, like for the ortho substituted phenols in Figure 1, the H-bond is disrupted upon an 180° rotation about the C–O bond.\nFigure 3 OCCN gauche structures with an intramolecular H-bond for 2-aminoethanol (7) and 2-NO2 ethanol (9); Conformations 8 and 10 indicate disrupted H-bonds after rotations by approximately 120° about the O–C axes. In aqueous solutions, the O (solute)…O (water) and N (solute)…O (water) radial distribution functions show their first minima at up to 350 pm [20,21]. This value has been accepted as the boundary of the first hydration shell around the polar sites of solutes. This, however, does not mean that intermolecular H-bonds would be expected with X (solute)…O (water) separation up to 350 pm in solution. Analyses of Monte Carlo results (see below) always point out that the number of the solvent molecules engaged in H-bond(s) to the solute is smaller than the total number of the solvent molecules in the first hydration shell(s) around the polar site(s). The solute-solvent pair-energy distribution functions show, in general, a maximum and a minimum within the range of −70 to about −10 kJ/mol for aqueous solutions. Integration of this distribution function up to its first minimum was interpreted by Jorgensen et al., [20] as the number of the intermolecular, solute-solvent hydrogen bonds in water.\nA recent review by Nagy [22] dealt with the in-solution conformational/tautomeric equilibria for small molecules in general, and the theoretical methods applied in the corresponding calculations were shortly characterized in that review. References to the theoretical methods will be only given for some less-known methods in the present paper. Basis sets applied in quantum mechanical calculations will be provided in cases where they may be needed to evaluate the relevance of the obtained results. Regarding structural analyses, recent publications were mostly sought with the hope that meaningful former studies would be listed in the later ones.\nAs stated in the title, the present survey emphasizes a special structural problem. Regarding the methodology, only problems related to the modeling of the H-bond will be discussed here. The conformational issue will be investigated for a number of families of small molecules. 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