The Origin of the Non-Additivity in Resonance-Assisted Hydrogen

Article Cite This: J. Phys. Chem. A XXXX, XXX, XXX-XXX

pubs.acs.org/JPCA

The Origin of the Non-Additivity in Resonance-Assisted Hydrogen Bond Systems Published as part of The Journal of Physical Chemistry virtual special issue “Manuel Yáñez and Otilia Mó Festschrift”. Xuhui Lin,† Huaiyu Zhang,† Xiaoyu Jiang,‡ Wei Wu,† and Yirong Mo*,§ †

The State Key Laboratory of Physical Chemistry of Solid Surfaces, iChEM, Fujian Provincial Key Laboratory of Theoretical and Computational Chemistry and College of Chemistry and Chemical Engineering, Xiamen University, Xiamen, Fujian 361005, China ‡ College of Ecological Environment and Urban Construction, Fujian University of Technology, Fuzhou 350108, China § Department of Chemistry, Western Michigan University, Kalamazoo, Michigan 49008, United States ABSTRACT: The concept of resonance-assisted hydrogen bond (RAHB) has been widely accepted, and its impact on structures and energetics can be best studied computationally using the block-localized wave function (BLW) method, which is a variant of ab initio valence bond (VB) theory and able to derive strictly electron-localized structures self-consistently. In this work, we use the BLW method to examine a few molecules that result from the merging of two malonaldehyde molecules. As each of these molecules contains two hydrogen bonds, these intramolecular hydrogen bonds may be cooperative or anticooperative, depended on their relative orientations, and compared with the hydrogen bond in malonaldehyde. Apart from quantitatively confirming the concept of RAHB, the comparison of the computations with and without π resonance shows that both σ-framework and πresonance contribute to the nonadditivity in these RAHB systems with multiple hydrogen bonds.

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acceptor through the π conjugation,30,31 and the quenching of the π resonance would significantly weaken the hydrogen bonding to a level similar to saturated counterparts. Furthermore, we showed that the intramolecular hydrogenbonding strength can be reflected by the classical electrostatic attraction between the H-bond donor and the acceptor whose charges can be easily approximated with population analyses.29 Differently, Góra et al. showed that the π resonance effect contributes little to the stability of RAHB systems based on another variant VB method called complete active space VB (CASVB).32,33 The difference between our BLW method and the CASVB method lies in the different definition of orbitals. The BLW method uses strictly localized orbitals, which expand only on molecular fragments (functional groups or individual atoms), while in the CASVB method, overlap-enhanced orbitals, which are expanded in the whole system thus delocalized just like MOs, are used. The use of delocalized orbitals makes any particular VB structure contaminated by all others and consequently underestimates the resonance stabilization. With the establishment that the conjugation moves partial π electrons from the H-bond donor to the acceptor and thus increases the negative charge on the latter as exemplified by malonaldehyde (1) in Scheme 1, for a system with more than one intramolecular RAHBs such as merged structures of

INTRODUCTION Since the proposal of the intramolecular resonance-assisted hydrogen bond (RAHB) concept,1−6 extensive studies have been conducted both experimentally7,8 and computationally,9−14 in attempts to understand and utilize the RAHB phenomenon, where the intramolecular H bond is strengthened with the shortening of bond, decreasing of the O−H vibrational frequency, and abnormal downfield 1H NMR chemical shift. RAHB highlights the interplay between resonance and hydrogen bonding and has been found to be crucial in protein folding and DNA pairing.15−18 Yet there are controversies over the nature of RAHB, as there is a possibility that the stronger intramolecular RAHB in unsaturated compounds than in their saturated analogues does not necessarily means that there is a RAHB phenomenon; rather, it may simply result from the constraints imposed by the σ-skeleton framework.19−21 But it may be inappropriate to compare a conjugated system with its saturated analogue, as they have quite different σ-skeleton frameworks. The most objective reference to analyze the σskeleton impact on RAHB is the use of the molecule itself but with π conjugation disabled. Most recently, we used the blocklocalized wave function (BLW) method,22−24 which is the simplest variant of ab initio valence bond (VB) theory25−28 but incorporates the efficiency of the molecular orbital (MO) theory as well, to study a series of paradigmatic examples of RAHBs and theoretically confirmed the existence of the RAHB phenomenon.29 In other words, the enhancement of the intramolecular hydrogen bonding in RAHB systems indeed results from the charge flow from the H-bond donor to the © XXXX American Chemical Society

Received: September 22, 2017 Revised: October 19, 2017 Published: October 19, 2017 A

DOI: 10.1021/acs.jpca.7b09425 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

where ML is the normalization constant, Â is the antisymmetrizer, and φ2i−1,2i is a bond function composed of nonorthogonal orbitals ϕ2i−1 and ϕ2i (or a lone pair if ϕ2i−1 = ϕ2i)

Scheme 1. Examples of Hydrogen Bond Influenced by π Conjugation

φ2i − 1,2i = Â {ϕ2i − 1ϕ2i[α(i)β(j) − β(i)α(j)]}

Since the HLSP can be expanded a number of Slater determinants, the computational costs rise swiftly with the increasing of bonding electron pairs. One way to remarkably increase the efficiency is to represent the bond function (eq 2) with a doubly occupied yet strictly localized group orbital. The BLW method further extends this idea by allowing the doubly occupied orbitals to expand within fragments of a molecule; that is, all electrons and primitive basis functions (χμ) can be partitioned into several subgroups (blocks), and each orbital is block-localized and expanded in only one block and thus called block-localized MO (BL-MO).22−24 Assuming that there are mi basis functions and ni electrons for block i, we express BL-MOs for this block as

malondialdehyde,34 it has been proposed that the π resonance may not be necessarily cooperative with the hydrogen bonding, and the two forces can be cooperative (2 in Scheme 1) or anticooperative (3 and 4).1,6,34−36 Yet we note that any hydrogen bonding can be affected by the presence of another H bond, whether there is π conjugation involved or not. In other words, the H bonds in 2 will be enhanced mutually and in 3 and 4 will be weakened each other even without the π conjugation, due to the antiparallel and parallel of the H-bonding directions. To clarify the role of π resonance in the H-bond strength in these systems, it is essential to find new approaches. Most recently, RomeroMontalvo et al. explored these systems with the Interacting Quantum Atoms (IQA) energy partition and the Quantum Theory of Atoms in Molecules (QTAIM) and found that the cooperativity and anticooperativity therein are related to a greater or lesser strengthening of the bicyclic structures of these systems.34 But these approaches are based on the total molecular electron density and thus inconvenient to differentiate the σ and π electrons. Cooperativity or anticooperativity may simply result from the relative orientation of two neighboring H bonds, as each H bond has a local dipole that is associated with the electron transfer (hyperconjugation). Considering the importance of hydrogen bonding in determining the three-dimensional structures adopted by proteins and nucleic bases, the strength of polymers, the pattern of self-assembling materials, and much more, we believe that it is essential to interpret the bonding in terms of various components and clarify the interplay between π resonance and the H-bond strength. The results can provide us not only deep insights into the nature of RAHB (or HB in general) but also guide the development of force fields that have been broadly used in the computational simulations of biological systems and materials and the rational design of novel materials assembled via hydrogen bonds. In this work, we employed the BLW to explore the interplay between π conjugation and the cooperativity/anticooperativity in molecules of multiple H bonds.

φji =

mi

∑ Cjiμχμi (3)

μ=1

The electron-localized state for a closed-shell subsequently can be defined using a Slater determinant as ΨBLW = det|(φ11)2 (φ21)2 ···(φn1 /2)2 ···(φ1i)2 ···(φni /2)2 ···(φnk /2)2 | 1

= Â [Φ1 ··· Φi ··· Φk ]

i

k

(4)

where Φi is the direct product of block-localized orbitals in block i. Like MO theory, we constrain the orbitals in the same subspace to be orthogonal, but orbitals belonging to different blocks are free to overlap and nonorthogonal. In this way, the BLW method combines the advantages and characteristics of both MO and VB theories. The BLW method is available at the DFT level with the geometry optimization and frequency computation capabilities.22,24 H-Bonding Strength Evaluation. Unlike intermolecular H bonds, which can be conveniently evaluated through a supramolecular approach, intramolecular hydrogen bonding cannot be accurately measured. The simplest and most effective approximation is to flip the H-bond donor, that is, the hydroxyl group in present cases 1−4, by 180° from its hydrogen-bonding position, and the subsequent energy variation is taken as the bonding strength.34,39−57 The comparison of the strengths computed with the regular DFT and BLW-DFT methods highlights the role of π resonance in hydrogen bonding. As there are two H bonds in the merged structures of malondialdehyde 2−4, flipping both hydroxyl groups breaks both H bonds. Thus, ideally we would like to compare the cooperativity or anticoopertivity in systems 2−4 with the bonding strength in malonaldehyde (1) using the following term

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METHODOLOGY Block-Localized Wave Function Method. Currently computational chemistry is dominated by methods based on MO theory, which assumes that all electrons delocalize to the whole studied system. With a different philosophy but the identical ultimate goal, VB theory starts from local atomic orbitals to build electron-localized Lewis (resonance) structures. When one resonance structure is not enough to describe the focused molecule, more resonance structures are introduced, as the resonance theory states.37,38 Within the VB theory, each resonance structure can be defined with a Heitler− London−Slater−Pauling (HLSP) function as ΨL = MLA(̂ φ1,2φ3,4 ··· φ2n − 1,2n)

(2)

flip flip ΔΔEcoop = ΔE2,3,4 − 2ΔE1flip

(5)

If ΔΔEcoop is larger than zero, the two H bonds are cooperative; otherwise, they are anticooperative. The comparison of the regular DFT (with π resonance on) and BLW-DFT (with π resonance off) thus reveals the nature of the cooperativity or anticooperativity in systems 2−4. However, we will show that this technique is problematic for systems with multiple H bonds, as when all H bonds are flipped, the hydroxyl groups

(1) B


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The Journal of Physical Chemistry A still can be viewed as local dipoles, and they will interact with each other. As such, the reference states are not comparable. Here we used the strategy adopted by Romero-Montalvo et al.34 and compared the flipping energy for one H bond with the other H bond formed (IN) or broken (OUT). As such, the cooperativity is measured as flip flip ΔΔEcoop = ΔE IN − ΔEOUT

(6)

If ΔΔEcoop is larger than zero, the two H bonds are cooperative; otherwise, they are anticooperative. The difference between this work and Romero-Montalvo et al.’s34 lies in the present BLW computations, which can measure the individual roles of π resonance and σ framework in cooperativity or anticooperativity. Computational Details. Throughout the work standard B3LYP DFT calculations with the basis set of 6-311+G(d,p) were performed, as this level of theory has been extensively applied to molecules with intramolecular hydrogen bonds and shown comparable with MP2/6-311+G(d,p) results.20,21 All computations were performed and geometries were optimized and justified by all real vibrational frequencies with the in-house version of the GAMESS software,58 where the BLW code was ported to. The comparison of the geometrical parameters and vibrational frequencies computed with the standard B3LYP and the BLW-DFT methods reveals the impact of π resonance on both the structures and energetics of molecules with intramolecular RAHBs.

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Figure 1. Structural parameters (in Å) in optimal geometries of 1−4 and their flipped 1′−4′ optimized at the B3LYP/6-311+G(d,p) level. The black values refer to regular DFT results, whereas the red ones are from the BLW method, which strictly localizes all π electron pairs. Data in italics and green are from partial BLW optimizations by restraining all main chain distances.

RESULTS AND DISCUSSION Influence of π Resonance on Geometries. We showed, with a series of molecules with one RAHB, that π resonance can significantly shorten the hydrogen bonds but stretch the bonds within the H-bond donor or acceptor.29 Here we similarly performed geometry optimizations for systems 2−4 and compared them with 1. Figure 1 shows the major structural parameters. For 2, the two H bonds (labeled a and e in Figure 1) are different. The O···H distance in the H bond e (1.637 Å) is much shorter than in malonaldehyde (1.703 Å), implying stronger bonding in the former H bond. This is accompanied by the stretching of the OH bond as well, but the carbonyl bond in the H bond e (1.228 Å) is a little shorter than in malonaldehyde (1.238 Å). In contrast, the O···H distance in the H bond a (1.745 Å) is longer than in malonaldehyde, implying the reduced bonding strength. Accordingly, both the OH and CO bonds in a are shorter than in malonaldehyde. If both hydroxyl groups are flipped out of the hydrogen bonding range, we observe a significant shortening of the OH bonds, so are the carbonyl bonds. However, if we retain the hydrogen bonding but strictly localize the π electrons, the hydrogen bonding distances will be notably stretched to 1.913 and 2.015 Å. Note that now the distance in the H bond e (2.015 Å) is longer (i.e., weaker) than in the H bond a (1.913 Å). In the electronlocalized state (red values in Figure 1), the optimal CC and CO bond lengths are very similar to the data in saturated molecules, confirming the rationality of the BLW method. The comparison of the DFT and BLW geometries reinforces our recent conclusion that π resonance can indeed enhance intramolecular H bonds.29,59 But still, with the π conjugation shut down, the H bonding is slightly stronger than in malonaldehyde as evidenced by the O···H distances.

The two H bonds in 3 are parallel and symmetrical, yet we find the shortest O···H distances (1.605 Å) here, implying the possibly strongest H-bonding strength. In accord, the carbonyl bond and the hydroxyl bond lengths reach 1.240 and 1.016 Å, respectively. Since the π resonance moves the electron density from the hydroxyl oxygen atoms to the carbonyl oxygen atoms through the same central CC double bond, interestingly, this CC double bond is greatly stretched to 1.426 Å. But once the conjugation is deactivated, it returns to 1.320 Å, comparable to the bond distance in ethylene. In fact, all structural parameters return to normal values if the π conjugation is quenched; for example, the single CC bond length is 1.550 Å, comparable to the value in ethane. The carbonyl bond and the hydroxyl bond lengths are 1.200 and 0.969 Å, respectively. 4 is an interesting molecule that has two H bonds sharing the same H-bond acceptor. Yet the H-bonding distances are comparable to or even slightly shorter than the H bond in malonaldehyde. As the only carbonyl group forms two H bonds, its CO bond is remarkably stretched to 1.275 Å, though both hydroxyl groups (0.996 Å) are normal and comparable to the one in malonaldehyde. Once all π electrons are strictly localized, the carbonyl CO bond length returns to 1.206 Å, comparable to the value in 1 (1.202 Å). Thus, the very long carbonyl bond length primarily results from the π conjugation, as both the hydroxyl groups lose π electron density to the carbonyl group through the central CC double bonds. The above full BLW optimizations show that the deactivation of π resonance not only stretches the intramolecular H bonds C


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Figure 2. EDD isosurface maps with the isovalue of 0.003 au showing the movement of electron density due to π conjugation. The orange and cyan colors refer to the increasing and decreasing of the electron density, respectively.

Table 1. Vibrational Frequency of the Hydroxyl Group (cm−1) and the Hydrogen Bonding Strength (kJ/mol) in ElectronDelocalized (DFT) and Electron-Localized (BLW) States, along with the Adiabatic Resonance Energy (kJ/mol) in the Former with Reference to the Latter molecule

method

vOH

ARE

ΔEflip

ΔΔEflip coop

ΔEflip IN

ΔEflip OUT

ΔΔEcoop

1

DFT BLW DFT BLW DFT BLW DFT BLW

3200 3747 3009/3312 3757/3696 2844/2884 3731/3732 3215/3248 3769/3774

158.6

54.1 27.8 84.0 35.4 144.4 56.4 120.2 58.8

0 0 −24.2 −20.2 36.2 0.8 12.0 3.2

26.9(64.3)a 10.1(30.1)a 21.8 8.5 56.9 26.8

19.7(57.1)a 5.3(25.3)a 122.6 47.8 63.3 32.0

7.2 4.8 −100.8 −39.3 −6.4 −5.2

2 3 4 a

323.3 365.1 313.2

Flipping energies for H bond 2a (2e) with the other 2e (2a) formed (IN) or broken (OUT).

Influence of π Resonance on Energetics. According to the original resonance theory, π resonance is a stabilizing force by definition, and its magnitude (called adiabatic resonance energy, or ARE) can be measured by the energy difference between the DFT and BLW (corresponding to the most stable resonance structure) optimal structures.37,38 There are six π electrons in malonaldehyde (1), and the merging of two malonaldehyde molecules leads to 2−4 with 10 π electrons. As listed in Table 1, resonance stabilizes 1 by 158.6 kJ/mol and 2− 4 by ∼2 times more. Of interest are 2 and 3, as they are isomers. 2 is marginally more stable than 3 by only 5.1 kJ/mol. Yet, Table 1 shows that π conjugation stabilizes 3 more than 2 by 41.8 kJ/mol. In other words, if there were no resonance, 2 would be more stable than 3 by 46.9 kJ/mol, which reflects the favorable orientations of two H bonds in 2. Once the hydroxyl groups are flipped out, the ARE values are 278.6 and 279.3 kJ/ mol for 2 and 3. Thus, π conjugation in 1−4 not only significantly varies structural parameters but also notably stabilizes molecules. These results are in accord with our conventional understanding of conjugation. Influence of π Resonance on the −OH Vibrational Frequency. It has been well-recognized that the formation of a H bond D−H···A leads to the elongation of the D−H bond and its accompanying red-shifting of the stretching vibrational

significantly but also alters the intramolecular CC and CO bond lengths remarkably. One subsequent question is whether the elongation of the H bonds is truly caused by the π electron density movement or simply a secondary effect from the variation of the geometries along the main chains. To address this question, we performed partial BLW optimizations by restraining all CC and CO bond distances at their DFT optimal values and examined the H-bond distances under constraints. Optimal OH bond and O···H distances from these partial BLW optimizations, as shown in Figure 1 in italics and green, are very close to the data from full BLW optimizations. Thus, the stretching of the H bonds indeed primarily results from the π electron density shifting. The movement of electron density due to the π conjugation can be better visually depicted through the electron density difference (EDD) maps, which is the difference between the DFT and BLW total electron densities. Figure 2 plots the EDD maps for systems 1−4. Apart from the loss (in cyan color) of πelectron densities on the hydroxyl groups and the gain (in orange color) of electron densities on the carbonyl groups, there are also the loss of electron density on CC double bonds and the gain of electron density on CC single bonds. Besides, the π conjugation also causes the counter polarization of the σ electron density, though the magnitude is small. D


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Figure 3. Electrostatic potentials on the 0.001 au electron density surfaces of molecules 1−4 computed with both the regular DFT and BLW methods at the B3LYP/6-311+G(d,p) theoretical level.

frequency.60−62 Using an empirical VB two-state model, Thompson and Hynes demonstrated the correlation between the frequency shifts and the charge transfer interaction.63 In general, the red shift with enhanced intensity in IR spectra is a “fingerprint” of hydrogen bonds, though it has been known that there are also blue-shifting phenomena.64−68 For the present cases, we observed significant enhancement of the red shifting of the hydroxyl groups due to the π conjugation (in the range of 400−900 cm−1 without scaling; see Table 1), in line with the definition of RAHB. For comparison, the OH vibrational frequency in methanol is 3842 cm−1 at the same theoretical level. In other words, if there were no π resonance in malonaldehyde (1), the red shifting would be only 95 cm−1. The π resonance further red-shifts the OH vibrational frequency by 547 cm−1. This is echoed by the change of the hydrogen bonding strength from 27.8 to 54.1 kJ/mol. More red shifting can be found in 2 and 3 due to π resonance. Correlation between Classical Electrostatic Attraction and H-Bond Strength. The flipping of the hydroxyl group in malonaldehyde leads to the energy increasing 54.1 kJ/mol. This is a reasonable amount for the H-bonding strength. With the π resonance deactivated, however, the flipping causes the energy change by 27.8 kJ/mol. In other words, half of the H-bond strength in malonaldehyde comes from the resonance enhancement. The existence of multiple H bonds in one molecule results in the coupling (nonadditivity) of these H bonds as studied by others, notably Romero-Montalvo et al.,1,6,34−36 who extensively explored the potential energy curves along the flipping of hydroxyl groups. By individually forming H bonds, Romero-Montalvo et al. showed that the formation energy (Hbonding strength) of H bond a (or e) in 2 is larger in the presence of H bond e (or a), suggesting the cooperative effect. In contrast, the formation energy of a H bond is smaller in the presence of another H bond in 3 and 4, suggesting the anticooperative effect.34 Here examine the energy changes by flipping H bonds in 2−4, with and without the resonance assistance, as listed in Table 1, to differentiate the σ and π roles. We first flip both H bonds and compare the energy changes in 2−4 with 1, as prescribed in eq 5. There are two notable findings from the examination of the flipping energies in 2−4. One is the remarkable reduction (by more than 50%) of the flipping energies with the shutdown of the π resonance. This echoes the significant geometrical changes discussed in the above and once again shows the influence of the π resonance on intramolecular hydrogen bonding. The other is, compared

with malonaldehyde, the two H bonds in 2 are anticooperative, while in 3 and 4 they are cooperative. This is against common sense and the results of Romero-Montalvo et al.34 If we look at the flipped geometries 2′−4′, the two hydroxyl groups in 2′ are in opposite directions, and thus the favorable local dipole− dipole electrostatic interactions would lower the energy of the reference molecule, leading to a much reduced flipped energy (i.e., underestimation of the bonding strength). In contrast, the two hydroxyl groups are in the same directions in 3′, and the unfavorable electrostatic repulsion pushes its molecular energy up. As a consequence, we observe a much exaggerated flipped energy for 3′. Compared to 3′, the two hydroxyl groups in 4′ are partially repulsive, and thus the extra flipped energy (eq 5) is positive but not as large as in 3′. As the total flipping energy (eq 5) is not a good measure of the cooperativity of intramolecular H bonds, we use eq 6 to examine the flipping energy for individual H bonds and study its correlation with π resonance. As expected, computations confirm the cooperative effect in 2 and the anticooperative effect in 3 and 4. While both the cooperativity in 2 and the anticooperativity in 4 are relatively insignificant, BLW computations show that these effects are dominated by the σframework but reinforced by the π resonance. Differently, 3 exhibits significant anticooperativity, which is mainly due to the π resonance (61%). But still, the σ-framework makes notable contribution (39%). To appreciate the contributions from both σ and π components to the nonadditivity in 2−4, we compared their electrostatic potential maps with and without π resonance in Figure 3. The most obvious finding from Figure 3 is that, when there were no π resonance, the charge separation in all systems would be more severe. In other words, π resonance helps balance the charge distribution over whole molecules, and it particularly shows the reduction of electron density on the hydroxyl oxygen and the increase of electron density on the carbonyl carbon, in accord with the EDD maps in Figure 2. This movement favors the electron transfer from the carbonyl oxygen to the hydroxyl hydrogen and consequently enhances the hydrogen bonding.

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CONCLUSION In this work, we examined the impact of π resonance on the hydrogen bonding strength in three β-enolones that result from the merging of two malonaldehyde molecules and were recently E


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The Journal of Physical Chemistry A studied by Romero-Montalvo and co-workers.34 The uniqueness of our study lies in the strict localization of π electrons, that is, the disabling of the π resonance, using the BLW method, which is a variant of ab initio VB theory.22−24 The comparison between the optimal geometries and energetics with the π resonance turned “on” and “off” confirms the concept of RAHB and shows the remarkable shortening and strengthening of concerned hydrogen bonds.1−6 With multiple hydrogen bonds existing in the same molecule, they may be cooperative or anticooperative, depending on their relative orientations. Our computations show that both the σ framework and π resonance make comparable contributions to the nonadditivity of these hydrogen bonds.

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(8) Rospenk, M.; Majewska, P.; Czarnik-Matusewicz, B.; Sobczyk, L. Polarized IR spectra of resonance assisted hydrogen bond (RAHB) in 2-hydroxyazobenzenes. Chem. Phys. 2006, 326, 458−464. (9) Alkorta, I.; Elguero, J.; Mó, O.; Yañez, M.; Del Bene, J. E. Are resonance-assisted hydrogen bonds ‘resonance assisted’? A theoretical NMR study. Chem. Phys. Lett. 2005, 411, 411−415. (10) Fuster, F.; Grabowski, S. J. Intramolecular hydrogen bonds: The QTAIM and ELF characteristics. J. Phys. Chem. A 2011, 115, 10078− 10086. (11) Trujillo, C.; Sánchez-Sanz, G.; Alkorta, I.; Elguero, J.; Mó, O.; Yáñez, M. Resonance assisted hydrogen bonds in open-chain and cyclic structures of malonaldehyde enol: A theoretical study. J. Mol. Struct. 2013, 1048, 138−151. (12) Beck, J. F.; Mo, Y. How resonance assists hydrogen bonding interactions: An energy decomposition analysis. J. Comput. Chem. 2007, 28, 455−466. (13) Guevara-Vela, J. M.; Romero-Montalvo, E.; Costales, A.; Pendás, Á . M.; Rocha-Rinza, T. The nature of resonance-assisted hydrogen bonds: a quantum chemical topology perspective. Phys. Chem. Chem. Phys. 2016, 18, 26383−26390. (14) Kurczab, R.; Mitoraj, M. P.; Michalak, A.; Ziegler, T. Theoretical analysis of the resonance assisted hydrogen bond based on the combined extended transition state method and natural orbitals for chemical valence scheme. J. Phys. Chem. A 2010, 114, 8581−8590. (15) Steiner, T. The hydrogen bond in the solid state. Angew. Chem., Int. Ed. 2002, 41, 48−76. (16) Fonseca Guerra, C.; Bickelhaupt, F. M.; Snijders, J. G.; Baerends, E. J. The nature of the hydrogen bond in DNA base pairs: The role of charge transfer and resonance assistance. Chem. - Eur. J. 1999, 5, 3581−3594. (17) Fonseca Guerra, C.; Bickelhaupt, F. M.; Snijders, J. G.; Baerends, E. J. Hydrogen bonding in DNA base pairs: Reconciliation of theory and experiment. J. Am. Chem. Soc. 2000, 122, 4117−4128. (18) Mo, Y. Probing the nature of hydrogen bonds in DNA base pairs. J. Mol. Model. 2006, 12, 665−672. (19) Sanz, P.; Mó, O.; Yáñez, M.; Elguero, J. Non-resonance-assisted hydrogen bonding in hydroxymethylene and aminomethylene cyclobutanones and cyclobutenones and their nitrogen counterparts. ChemPhysChem 2007, 8, 1950−1958. (20) Sanz, P.; Mó, O.; Yáñez, M.; Elguero, J. Resonance-assisted hydrogen bonds: A critical examination. Structure and stability of the enols of β-diketones and β-enaminones. J. Phys. Chem. A 2007, 111, 3585−3591. (21) Sanz, P.; Mó, O.; Yáñez, M.; Elguero, J. Bonding in tropolone, 2-aminotropone, and aminotroponimine: No evidence of resonanceassisted hydrogen-bond effects. Chem. - Eur. J. 2008, 14, 4225−4232. (22) Mo, Y. Geometrical optimization for strictly localized structures. J. Chem. Phys. 2003, 119, 1300−1306. (23) Mo, Y.; Peyerimhoff, S. D. Theoretical analysis of electronic delocalization. J. Chem. Phys. 1998, 109, 1687−1697. (24) Mo, Y.; Song, L.; Lin, Y. The block-localized wavefunction (BLW) method at the density functional theory (DFT) level. J. Phys. Chem. A 2007, 111, 8291−8301. (25) Wu, W.; Su, P.; Shaik, S.; Hiberty, P. C. Classical valence bond approach by modern methods. Chem. Rev. 2011, 111, 7557−7593. (26) Valence Bond Theory; Cooper, D. L., Ed.; Elsevier: Amsterdam, The Netherlands, 2002. (27) Gallup, G. A. Valence Bond Methods: Theory and Applications; Cambridge University Press: New York, 2002. (28) Shaik, S. S.; Hiberty, P. C. A Chemist’s Guide to Valence Bond Theory; Wiley: Hoboken, NJ, 2008. (29) Jiang, X.; Zhang, H.; Wu, W.; Mo, Y. A critical check for the role of resonance in intramolecular hydrogen bonding (IMHB). Chem. Eur. J. 2017, DOI: 10.1002/chem.201703952 (in press). (30) Desiraju, G. R.; Steiner, T. The Weak Hydrogen Bond In Structural Chemistry and Biology; Oxford University Press: New York, 2001. (31) Steiner, T. Structural evidence for resonance-assisted O−H···S hydrogen bonding. Chem. Commun. 1998, 3, 411−412.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Yirong Mo: 0000-0002-2994-7754 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS W.W. thanks the National Natural Science Foundation of China (No. 21290193) for support. Y.M. thanks the Faculty Research and Creative Activities Award, Western Michigan University, for support.

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REFERENCES

(1) Bertolasi, V.; Gilli, P.; Ferretti, V.; Gilli, G. Evidence for resonance-assisted hydrogen bonding. 2. Intercorrelation between crystal structure and spectroscopic parameters in eight intramolecularly hydrogen bonded 1,3-diaryl-1,3-propanedione enols. J. Am. Chem. Soc. 1991, 113, 4917−4925. (2) Gilli, G.; Bellucci, F.; Ferretti, V.; Bertolasi, V. Evidence for resonance-assisted hydrogen bonding from crystal-structure correlations on the enol form of the β-diketone fragment. J. Am. Chem. Soc. 1989, 111, 1023−1028. (3) Gilli, P.; Bertolasi, V.; Ferretti, V.; Gilli, G. Evidence for resonance-assisted hydrogen bonding. 4. Covalent nature of the strong homonuclear hydrogen bond. Study of the O-H–O system by crystal structure correlation methods. J. Am. Chem. Soc. 1994, 116, 909−915. (4) Gilli, P.; Bertolasi, V.; Ferretti, V.; Gilli, G. Evidence for intramolecular N-H···O resonance-assisted hydrogen bonding in βenaminones and related heterodienes. A combined crystal-structural, IR and NMR spectroscopic, and quantum-mechanical investigation. J. Am. Chem. Soc. 2000, 122, 10405−10417. (5) Gilli, P.; Bertolasi, V.; Pretto, L.; Ferretti, V.; Gilli, G. Covalent versus electrostatic nature of the strong hydrogen bond: Discrimination among single, double, and asymmetric single-Well hydrogen bonds by variable-temperature X-ray crystallographic methods in βdiketone enol RAHB systems. J. Am. Chem. Soc. 2004, 126, 3845− 3855. (6) Gilli, P.; Bertolasi, V.; Pretto, L.; Lyčka, A.; Gilli, G. The nature of solid-state N-H···O/O-H···N tautomeric competition in resonant systems. Intramolecular proton transfer in low-barrier hydrogen bonds formed by the ···OC-CN-NH···↔···HO-CC-NN··· ketohydrazone-azoenol system. A variable-temperature X-ray crystallographic and DFT computational study. J. Am. Chem. Soc. 2002, 124, 13554− 13567. (7) Chin, J.; Kim, D. C.; Kim, H. J.; Panosyan, F. B.; Kim, K. M. Chiral shift reagent for amino acids based on resonance-assisted hydrogen bonding. Org. Lett. 2004, 6, 2591−2593. F


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(53) González, L.; Mó, O.; Yáñez, M. High-level ab initio calculations on the intramolecular hydrogen bond in thiomalonaldehyde. J. Phys. Chem. A 1997, 101, 9710−9719. (54) Janusz Grabowski, S. J. π-Electron delocalisation for intramolecular resonance assisted hydrogen bonds. J. Phys. Org. Chem. 2003, 16, 797−802. (55) Zhang, Y.; Wang, C. S. Estimation on the intramolecular 10membered ring NH···OC hydrogen-bonding energies in glycine and alanine peptides. J. Comput. Chem. 2009, 30, 1251−1260. (56) Lenain, P.; Mandado, M.; Mosquera, R. A.; Bultinck, P. Interplay between hydrogen-bond formation and multicenter pelectron delocalization: Intramolecular hydrogen bonds. J. Phys. Chem. A 2008, 112, 10689−10696. (57) Nowroozi, A.; Hajiabadi, H. How to estimate the intramolecular hydrogen-bond energy of complex RAHB systems? A theoretical study. Struct. Chem. 2014, 25, 215−220. (58) Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S. T.; Gordon, M. S.; Jensen, J. J.; Koseki, S.; Matsunaga, N.; Nguyen, K. A.; Su, S.; Windus, T. L.; Dupuis, M.; Montgomery, J. A. J. General atomic and molecular electronic structure system. J. Comput. Chem. 1993, 14, 1347−1363. (59) Jiang, X.; Wu, W.; Mo, Y. The strength of intramolecular hydrogen bonds (IMHBs).Acta Phys.-Chim. Sin. 2017, doi: 10.3866/ PKU.WHXB201708174 (in press). (60) Ratajczak, H. Charge-transfer properties of the hydrogen bond. I. Theory of the enhancement of dipole moment of hydrogen-bonded systems. J. Phys. Chem. 1972, 76, 3000−2004. (61) Allerhand, A.; Von Rague Schleyer, P. v. R. A survey of C-H groups as proton donors in hydrogen bonding. J. Am. Chem. Soc. 1963, 85, 1715−1723. (62) Zierkiewicz, W.; Jurečka, P.; Hobza, P. On differences between hydrogen bonding and improper blue-shifting hydrogen bonding. ChemPhysChem 2005, 6, 609−617. (63) Thompson, W. H.; Hynes, J. T. Frequency shifts in the hydrogen-bonded OH stretch in halide-Water clusters. The importance of charge transfer. J. Am. Chem. Soc. 2000, 122, 6278− 6286. (64) Trudeau, G. T.; Dumas, J. M.; Dupuis, P.; Guerin, M.; Sandorfy, C. Intermolecular interactions and anesthesia: Infrared spectroscopic studies. Top. Curr. Chem. 1980, 93, 91−125. (65) Buděsí̌ nský, M.; Fiedler, P.; Arnold, Z. Triformylmethane: An efficient preparation, some derivatives, and spectra. Synthesis 1989, 1989, 858−860. (66) Hobza, P.; Špirko, V.; Havlas, Z.; Buchhold, K.; Reimann, B.; Barth, H. D.; Brutschy, B. Anti-hydrogen bond between chloroform and fluorobenzene. Chem. Phys. Lett. 1999, 299, 180−186. (67) Hobza, P.; Havlas, Z. Blue-shifting hydrogen bonds. Chem. Rev. 2000, 100, 4253−4264. (68) Mo, Y.; Wang, C.; Guan, L.; Braïda, B.; Hiberty, P. C.; Wu, W. On the nature of blue-shifting hydrogen bonds. Chem. - Eur. J. 2014, 20, 8444−8452.

(32) Cooper, D. L.; Thorsteinsson, T.; Gerratt, J. Modern VB representations of CASSCF wave functions and the fully-variational optimization of modern VB wave functions using the CASVB strategy. Adv. Quantum Chem. 1998, 32, 51−67. (33) Góra, R. W.; Maj, M.; Grabowski, S. J. Resonance-assisted hydrogen bonds revisited. Resonance stabilization vs. charge delocalization. Phys. Chem. Chem. Phys. 2013, 15, 2514−2522. (34) Romero-Montalvo, E.; Guevara-Vela, J. M.; Costales, A.; Pendás, Á . M.; Rocha-Rinza, T. Cooperative and anticooperative effects in resonance assisted hydrogen bonds in merged structures of malondialdehyde. Phys. Chem. Chem. Phys. 2017, 19, 97−107. (35) Bertolasi, V.; Pretto, L.; Gilli, G.; Gilli, P. π-Bond cooperativity and anticooperativity effects in resonance-assisted hydrogen bonds (RAHBs). Acta Crystallogr., Sect. B: Struct. Sci. 2006, 62, 850−863. (36) Otero, R.; Schöck, M.; Molina, L. M.; Lagsgaard, E.; Stensgaard, I.; Hammer, B.; Besenbacher, F. Guanine quartet networks stabilized by cooperative hydrogen bonds. Angew. Chem., Int. Ed. 2005, 44, 2270−2275. (37) Pauling, L. C. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, NY, 1960. (38) Wheland, G. W. The Theory of Resonance; John Wiley & Sons: New York, 1944. (39) Grabowski, S. J. Ab initio calculations on conventional and unconventional hydrogen bonds-study of the hydrogen bond strength. J. Phys. Chem. A 2001, 105, 10739−10746. (40) Grabowski, S. J. Hydrogen bonding strength - measures based on geometric and topological parameters. J. Phys. Org. Chem. 2004, 17, 18−31. (41) Wang, C. S.; Zhang, Y.; Gao, K.; Yang, Z. Z. A new scheme for determining the intramolecular seven-membered ring N−H···OC hydrogen-bonding energies of glycine and alanine peptides. J. Chem. Phys. 2005, 123, 024307. (42) Jabłoński, M.; Kaczmarek, A.; Sadlej, A. J. Estimates of the energy of intramolecular hydrogen bonds. J. Phys. Chem. A 2006, 110, 10890−10898. (43) Liu, T.; Li, H.; Huang, M. B.; Duan, Y.; Wang, Z. X. Two-way effects between hydrogen bond and intramolecular resonance effect: An ab initio study on complexes of formamide and its derivatives with water. J. Phys. Chem. A 2008, 112, 5436−5447. (44) Deshmukh, M. M.; Gadre, S. R. Estimation of N−H···OC intramolecular hydrogen bond energy in polypeptides. J. Phys. Chem. A 2009, 113, 7927−7932. (45) Gilli, P.; Pretto, L.; Bertolasi, V.; Gilli, G. Predicting hydrogenbond strengths from acid-base molecular properties. The pK(a) slide Rule: Toward the solution of a long-lasting problem. Acc. Chem. Res. 2009, 42, 33−44. (46) Wendler, K.; Thar, J.; Zahn, S.; Kirchner, B. Estimating the hydrogen bond energy. J. Phys. Chem. A 2010, 114, 9529−9536. (47) Rozas, I. On the nature of hydrogen bonds: an overview on computational studies and a word about patterns. Phys. Chem. Chem. Phys. 2007, 9, 2782−2790. (48) Estácio, S. G.; Cabral do Couto, P.; Costa Cabral, B. J.; Minas da Piedade, M. E.; Martinho Simões, J. A. Energetics of intramolecular hydrogen bonding in di-substituted benzenes by the ortho−para method. J. Phys. Chem. A 2004, 108, 10834−10843. (49) Lipkowski, P.; Koll, A.; Karpfen, A.; Wolschann, P. An approach to estimate the energy of the intramolecular hydrogen bond. Chem. Phys. Lett. 2002, 360, 256−263. (50) Woodford, J. N. Density functional theory and atoms-inmolecules investigation of intramolecular hydrogen bonding in derivatives of malonaldehyde and implications for resonance-assisted hydrogen bonding. J. Phys. Chem. A 2007, 111, 8519−8530. (51) Latajka, Z.; Scheiner, S. Proton transfer in the ground and first excited triplet states of malonaldehyde. J. Phys. Chem. 1992, 96, 9764− 9767. (52) Scheiner, S.; Kar, T.; Cuma, M. Excited state intramolecular proton transfer in anionic analogues of malonaldehyde. J. Phys. Chem. A 1997, 101, 5901−5909. G


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