Kinetic Energy Density as a Predictor of Hydrogen-Bonded OH

Article pubs.acs.org/JPCA

Kinetic Energy Density as a Predictor of Hydrogen-Bonded OHStretching Frequencies Published as part of The Journal of Physical Chemistry virtual special issue “Veronica Vaida Festschrift”. Joseph R. Lane,*,† Anne S. Hansen,‡ Kasper Mackeprang,‡ and Henrik G. Kjaergaard‡ †

School of Science, University of Waikato, Private Bag 3105, Hamilton 3240, New Zealand Department of Chemistry, University of Copenhagen, Universitetsparken 5, DK-2100 Copenhagen, Denmark

‡

S Supporting Information *

ABSTRACT: This work considers the nature of the intermolecular hydrogen bond in a series of 15 different complexes with OH donor groups and N, O, P, or S acceptor atoms. To complement the existing literature, room-temperature gas-phase vibrational spectra of the methanol−pyridine, ethanol−pyridine, and 2,2,2-trifluoroethanol−pyridine complexes were recorded. These complexes were chosen, as they exhibit hydrogen bonds of intermediate strength as compared to previous investigations that involved strong or weak hydrogen bonds. Non Covalent Interactions (NCI) theory was used to calculate various properties of the intermolecular hydrogen bonds, which were compared to the experimental OH-stretching vibrational red shifts. We find that the experimental OH-stretching red shifts correlate strongly with the kinetic energy density integrated within the reduced density gradient volume that describes a hydrogen bond [G(s0.5)]. Given that vibrational red shifts are commonly used as a metric of the strength of a hydrogen bond, this suggests that G(s0.5) could be used as a predictor of hydrogen bonding strength. cm−1.10−13,15 Consequently, fewer investigations of these hydrogen bonds exist compared to those of the OH−N hydrogen bond, and at room temperature in the gas phase only limited studies are available.10,16−27 In the present work, we consider three examples of OH−N hydrogen bonds in complexes that involve pyridine, which is a weaker N-based acceptor molecule than DMA or TMA. It is expected that this will result in OH−N hydrogen bonds of intermediate strength, complementing the existing literature for strong OH−N hydrogen bonds,9,14,28 and for the weaker OH− O, OH−P, and OH−S hydrogen bonds.10−13 Three different OH donor molecules are considered, and we recorded gasphase room-temperature vibrational spectra of the methanol− pyridine, ethanol−pyridine, and 2,2,2-trifluoroethanol−pyridine complexes. Collectively, this provides gas-phase room-temperature OH-stretching vibrational spectra for a total of 15 different complexes recorded under similar if not identical experimental conditions. NCI theory has been applied to all 15 complexes to theoretically characterize the nature of the hydrogen bonds. NCI theory can be used to visualize the position and nature of noncovalent interactions by identifying regions in three-

1. INTRODUCTION The hydrogen bond (XH−Y) is a ubiquitous noncovalent interaction that is crucial in both biological and chemical processes.1−3 A variety of hydrogen bonds are possible, with the most well-known examples including X = F, N, O, and C and Y  N, O, and S, where the hydrogen-bond donor and acceptor molecules can be both neutral or charged compounds.4 The extensive literature on the hydrogen bond has enabled investigations of how it is affected when small and/or large perturbations are performed, for example, changing the hydrogen bond acceptor atom along the periods or down the groups in the periodic table. Much work has been done to try to understand how the strength of a hydrogen bond correlates with physical parameters, including various structural and spectroscopic properties, most commonly in solid or liquid phases.5−8 In the present work, we consider OH−Y (where Y  N, O, P, S) hydrogen bonds, which have been the focus of a number of gas-phase room-temperature vibrational spectroscopy studies by some of the authors over the past decade.9−13 The OH−N hydrogen bond, the strongest of the four, was first detected in the 1960s in the gas phase at room temperature, with large OH-stretching red shifts of ∼300 cm−1 for the complexes between methanol (MeOH) and dimethylamine (DMA)/trimethylamine (TMA).9,12,14 In contrast, the OH−O, OH−P, and OH−S hydrogen bonds are similar and exhibit more modest OH-stretching red shifts of 103−140 © 2017 American Chemical Society

Received: March 17, 2017 Revised: April 21, 2017 Published: April 24, 2017 3452

DOI: 10.1021/acs.jpca.7b02523 J. Phys. Chem. A 2017, 121, 3452−3460

Article

The Journal of Physical Chemistry A

Harmonic vibrational frequencies were also calculated to confirm that the optimized structures were indeed minima. Single-point energy calculations were completed for the ωB97X-D/aug-cc-pVTZ structures, using explicitly correlated coupled cluster theory [CCSD(T)-F12a] and the cc-pVDZ-F12 orbital basis set in MOLPRO2012.1.35 The CCSD(T)-F12a/ cc-pVDZ-F12 electron densities were exported from MOLPRO in MOLDEN format and subsequently converted to wave function format using the Molden2AIM program.36 The intermolecular hydrogen bonding interactions were investigated using NCI theory using the NCIPLOT and NCImilano programs, respectively.37,38 AIM theory was also used as a point of comparison, with the identification and properties of the bond critical points evaluated using the Multiwfn program.39

dimensional space where the reduced density gradient (s) and electron density [ρ(r)] are low. 29 NCI theory is a complementary approach to the more traditional Atoms in Molecules (AIM) topological analysis of electron density,30 and it offers some advantages particularly for weak intramolecular hydrogen-bonding interactions.31,32 To date, NCI theory has largely been used for a qualitative interpretation of hydrogen bond strength, although Contreras-Garcia et al. showed that the integrated volume of s is related to the calculated intermolecular potential for simple hydrogen-bonded complexes.33 In this work, we compare experimental OH-stretching red shifts for a series of OH−Y complexes with their calculated NCI properties; the magnitude of the former is commonly considered an experimental measure of hydrogen bond strength.4 This investigation expands on recent work by some of the authors, who showed that the OH-stretching frequency of substituted amino alcohols that exhibit an intramolecular OH−N hydrogen bond correlates well with the kinetic energy integrated within an NCI isosurface volume.32

4. RESULTS AND DISCUSSION A total of 15 different complexes are considered in this work. For 12 of these, the gas-phase IR spectra have been previously reported by some of the authors, and the reader is directed to these separate publications for further details.9−13,15,40,41 Vibrational spectra of the remaining three complexes, namely, MeOH−pyridine, EtOH−pyridine, and TFE−pyridine, are presented in Section 4.1 along with theoretical analysis of the intermolecular interactions using NCI theory. In Section 4.2, we compare the experimentally observed OH-stretching red shifts of all 15 complexes to various theoretically determined properties thought to be predictors of hydrogen bond strength. For clarity, in this context we consider the magnitude of the OH-stretching red shift to represent the strength of a hydrogen bond (vide infra).4 For the EtOH complexes, we present calculated properties for the gauche conformer. In the spectra of the complexes it is not possible to distinguish between the gauche and trans conformers, so we use the lower-energy gauche conformer. As shown in the Supporting Information, we find that the calculated NCI properties for the two conformers are very similar, typically differing by ∼2%. 4.1. Hydrogen Bonding in Pyridine Complexes. In Figure 1, we show spectra of the OH-stretching vibrations in the MeOH−pyridine, EtOH−pyridine, and TFE−pyridine

2. EXPERIMENTAL DETAILS Methanol (MeOH; Aldrich 154903, ≥99.9%), ethanol (EtOH; Kemetyl anhydrous, 99.9%), 2,2,2-trifluoroethanol (TFE; Aldrich T63002, ≥99%), and pyridine (Aldrich 270970 anhydrous, 99.8%) were purified by freeze, pump, and thaw cycles. The gas-phase IR spectra were recorded at room temperature with a VERTEX 70 (Bruker) FTIR spectrometer, fitted with a CaF2 beamsplitter, an MCT detector, and an MIR light source. Spectra were recorded with a 1 cm−1 resolution, 500 scans, and multireflection cells (Infrared Analysis, Inc.) equipped with KCl windows and Au-coated mirrors. Optical path lengths of 2.4 m (MeOH−pyridine) and 16 m (EtOH− pyridine and TFE−pyridine) were used. A glass vacuum line (J. Young, base pressure from 5 × 10−5 to 1 × 10−4 torr) was used for sample preparation. The samples were connected to the glass vacuum line using 3/8 in. vacuum fittings (Swagelok). Sample and base pressures were measured with Varian Pirani capacitance diaphragm (4 × 10−5 to 1125 torr, PCG750) and Agilent Technologies capacitance diaphragm (1 × 10−3 to 10 torr, CDG500) pressure gauges connected to the vacuum line. Spectra of the gas-phase complexes were obtained by subtracting the individual monomer spectra from the spectrum of the mixture (e.g., MeOH + pyridine).11 However, the monomer spectra were recorded at a slightly different pressure relative to the pressure of each monomer in the mixture. The monomer spectra were scaled and then subtracted from the mixture, and an accurate subtraction was found when a straight baseline was obtained in the regions of monomer absorbance. The individual monomer pressures in the mixture were obtained from the spectral subtraction, by multiplying the scaling factor and the monomer pressure from the individual monomer spectra. Spectral subtraction and analyses were performed with OPUS 6.5 (Bruker) and OriginPro 9.1. 3. COMPUTATIONAL DETAILS The structures of the complexes and their corresponding monomers were optimized with density functional theory (DFT) using the ωB97X-D functional with the aug-cc-pVTZ basis set. All DFT calculations were completed using Gaussian 09,34 with the opt = tight and int = UltraFine keywords to ensure that the optimized structures were accurately converged.

Figure 1. OH-stretching vibrations in the MeOH−pyridine, EtOH− pyridine, and TFE−pyridine complexes, respectively. The spectra were recorded with monomer pressures of 34 torr MeOH + 8.7 torr pyridine, 11 torr EtOH + 3.8 torr pyridine, and 5.7 torr TFE + 0.64 torr pyridine. 3453


Article


and the π-electrons of pyridine. For the TFE−pyridine complex, the volume between the hydrogen atom of the OH group and the nitrogen atom has a darker blue color than that for the MeOH and EtOH complexes, indicating that the TFE− pyridine complex exhibits a much stronger intermolecular hydrogen bond. There are also two distinct green volumes between the monomers that are best described as weak van der Waals interactions. The interaction between the hydrogen atom of pyridine and the fluorine atom of TFE is weakly attractive and best thought of as a very weak CH−F hydrogen bond, whereas the volume at the center of the pseudoring between the OH−N and CH−F interactions is weakly repulsive. In Figure 3 we present plots of the reduced density gradient (s) and sign(λ2)ρ for the MeOH−pyridine, EtOH−pyridine, and TFE−pyridine complexes to consider the nature of the intermolecular interactions more quantitatively. The strength of interaction is related to the value of ρ, with troughs at very negative values of sign(λ2)ρ corresponding to very attractive interactions and troughs at very positive values of sign(λ2)ρ corresponding to very repulsive interactions; that is, the larger the density (further from zero), the stronger the interaction.29,37 It is also worth noting the correspondence between the reduced density gradient and topological analysis of the electron density, as in the widely used AIM approach.30 Troughs where the minimum of s is equal to 0 correspond to critical points in the electron density that can be further categorized as either (3,−1) bond critical points [where sign(λ2)ρ is negative] or (3,+1) ring critical points [where sign(λ2)ρ is positive]. For all three complexes, the trough with the most negative value of sign(λ2)ρ corresponds to the intermolecular hydrogen bond between the hydrogen atom of the OH group and the nitrogen atom of pyridine. For MeOH−pyridine and EtOH− pyridine, the density at this bond critical point is 0.030 au indicating that the two hydrogen bonds are of similar strength. For TFE−pyridine, a stronger hydrogen bond is evidenced by a density of 0.038 au at the bond critical point. The secondary

complexes. Additional spectra of the complexes, recorded with different combinations of monomer pressures, are shown in Figures S1−S6 of the Supporting Information. In Table 1, we Table 1. Observed OH-Stretching Frequencies (cm−1) and Red Shifts (Δν̃, cm−1) for the MeOH−Pyridine, EtOH− Pyridine, and TFE−Pyridine Complexes ν̃

MeOH−pyridine

EtOH−pyridine

TFE−pyridine

monomer complex Δν̃a

3681 3472 209

3669 3462 207

3657 3289 368

a

Difference between the OH-stretching frequency of the monomer and complex. For the monomers MeOH and TFE we use the Qbranch, and for EtOH we use a weighted band center, due to the two conformers present.

summarize the corresponding frequencies and frequency shifts relative to the OH-stretching vibration in the monomers. As the acidity of alcohol hydrogen bond donor increases from MeOH/ EtOH to TFE,42,43 the OH-stretching frequency shift increases in agreement with previous observations.12,16,44 In Figure 2 we present NCI isosurfaces for the MeOH− pyridine, EtOH−pyridine, and TFE−pyridine complexes to illustrate the nature of the intermolecular hydrogen bond. The NCI isosurfaces shown in Figure 2 were obtained with CCSD(T)-F12a/cc-pVDZ-F12 electron densities using the corresponding ωB97X-D/aug-cc-pVTZ optimized structures. We have previously shown that this theoretical approach yields electron densities in excellent agreement to those obtained with fully optimized CCSD(T)-F12a/cc-pVDZ-F12 structures.32 The NCI isosurface of the MeOH−pyridine complex exhibits a single circular pale blue volume between the hydrogen atom of the OH group and the nitrogen atom of pyridine that corresponds to the intermolecular hydrogen bond. The EtOH− pyridine complex has an almost identical volume corresponding to the intermolecular hydrogen bond but also exhibits secondary interactions between the methyl CH of ethanol

Figure 2. NCI isosurfaces for MeOH−pyridine (top-left), EtOH−pyridine (top-right) and TFE−pyridine (bottom) obtained with the CCSD(T)F12a/cc-pVDZ-F12 method using the ωB97X-D/aug-cc-pVTZ optimized geometry. s = 0.5 au and a blue-green-red color scale from −0.06 < sign(λ2)ρ < +0.06. The atoms are colored as white (H), gray (C), blue (N), red (O), and green (F). 3454


Article


Figure 3. Plots of the reduced density gradient s and sign(λ2)ρ for MeOH−pyridine (top-left), EtOH−pyridine (top-right), and TFE−pyridine (bottom) obtained with the CCSD(T)-F12a/cc-pVDZ-F12 method using the ωB97X-D/aug-cc-pVTZ optimized geometry. The blue-green-red color scale is consistent with that used in Figure 2.

setup.9,10,12,13 Also included in Table 2 are the calculated CCSD(T)-F12a/cc-pVDZ-F12 interaction energies that were obtained using the ωB97X-D/aug-cc-pVTZ optimized geometries. These interaction energies are raw differences in electronic energies and are not corrected for zero-point vibrational energy or basis set superposition error. A summary of the optimized geometric parameters that describe the intermolecular hydrogen bonds can be found in Table S1 of the Supporting Information. As can be seen in Table 2 and Figure 4, a range of hydrogen bond strengths are considered in this investigation, with interaction energies spanning from 21.1 to 47.9 kJ mol−1 and OH-stretching red shifts from 103 to 488 cm−1. A correlation plot between the experimental OH-stretching red shifts and the calculated interaction energies is shown in Figure 4. We find a reasonable linear correlation between Δν̃ and the calculated interaction energies, with R2 = 0.82. The remaining scatter in the data is largely attributed to the presence of secondary interactions other than the intended OH−Y hydrogen bond, which will affect the interaction energy but not the OHstretching red shift. For example, the CH−F interaction shown in Figure 2 between the closest hydrogen atom of pyridine and fluorine atom of TFE is weakly attractive, resulting in a larger calculated interaction energy than would be otherwise expected given the experimentally observed OH-stretching red shift; that is, this leads to a data point below the linear trend line, which will decrease the slope. Consequently we chose to use the OHstretching red shift, rather than the calculated interaction

intermolecular interactions that are present in the EtOH− pyridine and TFE−pyridine complexes are also evident in the plots of s versus sign(λ2)ρ as the two troughs around zero. These are missing in the MeOH−pyridine plots. Finally, all three complexes exhibit a trough at positive values of sign(λ2)ρ that corresponds to the repulsive interaction at the center of the pyridine ring associated with ring closure. 4.2. Predictors of OH-Stretching Red Shifts. In Table 2 we present the experimentally observed OH-stretching red shifts associated with formation of the complexes. The corresponding vibrational spectra were all recorded under similar conditions, most with the same experimental Table 2. Experimental OH-Stretching Red Shifts (Δν̃ in cm−1) and Calculated CCSD(T)-F12a/cc-pVDZ-F12 Interaction Energies (Int. E. in kJ mol−1)a MeOH acceptor (N) pyridine (N) TMA (O) DME (P) TMP (S) DMS

Δν̃ b

209 326c 103d 140f 113d

Int. E. 30.7 34.5 21.1 24.8 23.7

EtOH Δν̃ b

207 319c 104e 138f 112e

TFE

Int. E. 31.7 35.4 22.5 25.9 25.2

Δν̃ b

368 488c 198e 267f 201e

Int. E. 47.9 43.8 35.4 30.3 32.3

a

All frequency shifts are taken as the observed band maxima of the complex minus the monomer frequencies in Table 1. bThis work. c Reference 12. dReference 10. eReference 15. fReference 13. 3455


Article


Figure 4. Correlation plot of experimental Δν̃ vs calculated interaction energies. Calculated results are obtained with the CCSD(T)-F12a/cc-pVDZF12 method using the ωB97X-D/aug-cc-pVTZ optimized geometries.

Figure 5. Correlation plot of Δν̃ vs ρ at smin. (top) Acceptor atom type is ignored (red diamonds). (bottom) Data are grouped by the acceptor atom type (colored diamonds). Calculated results are obtained with the CCSD(T)-F12a/cc-pVDZ-F12 method using the ωB97X-D/aug-cc-pVTZ optimized geometries.

complexes (top, shown as red diamonds) we find a reasonably linear correlation between Δν̃ and ρ at smin, with an overall R2 = 0.77. However, if we consider each of the acceptor atom types separately (bottom, shown as colored diamonds), then the correlation between Δν̃ and ρ at smin improves significantly. For the N acceptors, the R2 value becomes 0.99 (six data points), and while we only have three data points for the other acceptors making R2 values largely meaningless, there appears to be a better linear correlation between Δν̃ and ρ at smin.

energies, to describe the strength of the localized hydrogen bond, as this is largely unaffected by the presence of secondary interactions. As discussed in Section 4.1, the value of the density corresponding to the bottom of the attractive trough in s (equivalent to the density at the bond critical point in AIM theory) can be used to quantify the strength of the hydrogen bond interaction.29,37,45−47 In Figure 5, we plot the OHstretching red shift against ρ at smin. If we ignore the nature of the acceptor atom type and simply correlate the data for all 15 3456


Article


Figure 6. Correlation plot of Δν̃ vs the integrated volume of s0.5. Data points are grouped by the hydrogen bond donor molecule (top, open circles) and by the hydrogen bond acceptor (bottom, diamonds). Calculated results are obtained with the CCSD(T)-F12a/cc-pVDZ-F12 method using the ωB97X-D/aug-cc-pVTZ optimized geometries.

Figure 7. Correlation plot of Δν̃ vs G(s0.5). Calculated results are obtained with the CCSD(T)-F12a/cc-pVDZ-F12 method using the ωB97X-D/ aug-cc-pVTZ optimized geometries.

correlation (R2 = 0.0008) between Δν̃ and s0.5 if the nature of the acceptor atom is ignored (top, red diamonds). In contrast, if we consider each acceptor atom type separately, then there appears to be a strong correlation, with the N acceptors yielding a correlation coefficient of R2 = 0.94 (six data points) (bottom, blue diamonds). In Figure 7, we plot the OH-stretching red shift against the kinetic energy density integrated within the reduced density gradient volume with s = 0.5 a.u., G(s0.5). Saleh et al. have

In Figure 6, we plot the OH-stretching red shift against the integrated volume of the reduced density gradient within the s = 0.5 a.u. isosurface, s0.5. The integrated volume of s has been previously shown to be a well-behaved function of the intermolecular distance between simple hydrogen-bonded monomers, and hence it can be related to the intermolecular potential.33 However, it appears that the relationship between s0.5 and the strength of interaction is strongly dependent on the acceptor atom type. We find there to be essentially no 3457


Article


Figure 8. Correlation plot of Δν̃ vs G(r) at smin. (top) Acceptor atom type is ignored (red diamonds). (bottom) Data are grouped by the acceptor atom type (colored diamonds). Calculated results are obtained with the CCSD(T)-F12a/cc-pVDZ-F12 method using the ωB97X-D/aug-cc-pVTZ optimized geometries.

previously shown that G(s0.5) correlates well with the calculated interaction energy of a series of simple dimers.48 Very recently, some of the authors showed that experimentally measured OHstretching frequencies of substituted aminoalcohols that exhibit OH−N intramolecular hydrogen bonding interactions also exhibit a linear correlation to G(s0.5).32 A strong linear correlation is evident in Figure 7 between Δν̃ and G(s0.5) with R2 = 0.99 (Δν̃ = 55 750G(s0.5) − 164). This is an important result, linking an accepted experimental measure of hydrogen bond strength to a readily calculable property. The linear correlation between Δν̃ and G(s0.5) improves slightly if the results are grouped by either the donor molecule or by the acceptor molecule (Supporting Information, Figure S8). However, this improvement is not significant for the modest number of data points available and furthermore yields fitted equations that are more specific and cannot therefore be applied as generally to other systems. In Figure 8, we plot the OH-stretching red shift against the kinetic energy density G(r) evaluated at the minimum of smin (equivalent to the bond critical point in AIM theory). This is to assess whether it is necessary to integrate G(r) within s0.5 or if simply the value of G(r) at smin can be used as an effective predictor of Δν̃. If we ignore the nature of the acceptor atom type and correlate the data for all 15 complexes (top, shown as red diamonds) we find there to be a weak linear correlation between Δν̃ and G(r), with an overall R2 = 0.42. The correlation improves significantly if we consider each of the acceptor atom types separately (bottom, shown as colored

diamonds), similar to what was observed in Figure 5 for the relationship between Δν̃ and ρ at smin.

5. CONCLUSIONS We have investigated the nature of the hydrogen bonding interactions in a series of different complexes with OH donor groups and N, O, P, or S acceptor atoms. Specifically, we consider complexes with MeOH, EtOH, and TFE acting as the donor molecule and pyridine, TMA, DME, DMS, and TMP acting as the acceptor molecule. Room-temperature gas-phase FTIR spectra of the MeOH−pyrdine, EtOH−pyridine, and TFE−pyridine complexes are reported for the first time and compared to literature spectra of the other complexes, which were recorded under similar conditions.9−13 Complexes with a reasonably wide range of hydrogen bond strengths were considered, with experimentally observed OH-stretching red shifts from 103 to 488 cm−1 and calculated interaction energies spanning from 21.1 to 47.9 kJ mol−1. NCI theory was applied to the complexes using CCSD(T)F12a/cc-pVDZ-F12 electron densities calculated with the ωB97X-D/aug-cc-pVTZ optimized structures. Various comparisons were made between the experimental OH-stretching red shifts (Δν̃) and the calculated NCI properties to ascertain the best predictor of hydrogen bond strength. We showed that Δν̃ demonstrates a reasonable linear correlation with the value of the density (ρ) at the minimum of s (equivalent to the density at the bond critical point in AIM theory), which is improved significantly if the nature of the acceptor molecule is taken into consideration; that is, that each acceptor molecule needs to be 3458



■

fitted separately for a good correlation. Similarly, the integrated volume of the reduced density gradient (s0.5) and the kinetic energy [G(r)] at smin are only good predictors of the OHstretching frequency if the acceptor atom type is considered. In contrast, Δν̃ was found to give a strong linear correlation with the kinetic energy density integrated within the reduced density gradient volume [G(s0.5)], which is largely independent of the hydrogen bond acceptor atom type. We suggest that G(s0.5) could be used to estimate the strength of a hydrogen bond, which is of particular value for intramolecular hydrogen bonds where determining a zeroth-order wave function for the donor and acceptor fragments is nontrivial but is being actively worked on.49,50 We believe that G(s0.5) offers some advantages over the more conventional approach of analyzing electron density properties at bond critical points.30,45−47 Namely, that it can be applied to hydrogen-bonding interactions that do not exhibit bond critical points, such as constrained intramolecular hydrogen bonds,32,51−53 and that the interpretation of bond strength is more universal for different hydrogen bond acceptors33 (as shown in Figure 7). While the present results are encouraging, caution should be exercised if attempting to extrapolate the linear relationships beyond the range of data that is considered in this work. Future investigations should therefore include complexes that exhibit very weak and/or very strong OH−Y hydrogen bonds, although these are likely to be much more experimentally challenging in the gas phase. To facilitate a more realistic comparison between theory and experiment, an ensemble of different conformers should also be considered when calculating the NCI properties to simulate the effects of thermal fluctuations.54,55

■

REFERENCES

(1) Kulmala, M. How Particles Nucleate and Grow. Science 2003, 302, 1000−1001. (2) Kjaergaard, H. G.; Robinson, T. W.; Howard, D. L.; Daniel, J. S.; Headrick, J. E.; Vaida, V. Complexes of Importance to the Absorption of Solar Radiation. J. Phys. Chem. A 2003, 107, 10680−10686. (3) Schuster, P.; Wolschann, P. Hydrogen Bonding: From Small Clusters to Biopolymers. Monatsh. Chem. 1999, 130, 947−960. (4) Arunan, E.; Desiraju, G. R.; Klein, R. A.; Sadlej, J.; Scheiner, S.; Alkorta, I.; Clary, D. C.; Crabtree, R. H.; Dannenberg, J. J.; Hobza, P.; et al. Definition of the Hydrogen Bond (IUPAC Recommendations 2011). Pure Appl. Chem. 2011, 83, 1637−1641. (5) Novak, A. Hydrogen Bonding in Solids Correlation of Spectroscopic and Crystallographic Data. In Large Molecules; Structure and Bonding Series; Springer: Berlin, Germany, 1974; Vol. 18, pp 177−216. DOI: 10.1007/BFb0116438. (6) Mikenda, W. Stretching Frequency Versus Bond Distance Correlation of OD(H)−Y (Y  N, O, S, Se, Cl, Br, I) Hydrogen Bonds in Solid Hydrates. J. Mol. Struct. 1986, 147, 1−15. (7) Iogansen, A. Direct Proportionality of the Hydrogen Bonding Energy and the Intensification of the Stretching ν(XH) Vibration in Infrared Spectra. Spectrochim. Acta, Part A 1999, 55, 1585−1612. (8) Rozenberg, M.; Loewenschuss, A.; Marcus, Y. An Empirical Correlation Between Stretching Vibration Redshift and Hydrogen Bond Length. Phys. Chem. Chem. Phys. 2000, 2, 2699−2702. (9) Howard, D. L.; Kjaergaard, H. G. Vapor Phase Near Infrared Spectroscopy of the Hydrogen Bonded Methanol-Trimethylamine Complex. J. Phys. Chem. A 2006, 110, 9597−9601. (10) Howard, D. L.; Kjaergaard, H. G. Hydrogen Bonding to Divalent Sulfur. Phys. Chem. Chem. Phys. 2008, 10, 4113−4118. (11) Du, L.; Lane, J. R.; Kjaergaard, H. G. Identification of the Dimethylamine-Trimethylamine Complex in the Gas Phase. J. Chem. Phys. 2012, 136, 184305. (12) Hansen, A. S.; Du, L.; Kjaergaard, H. G. The Effect of Fluorine Substitution in Alcohol-Amine Complexes. Phys. Chem. Chem. Phys. 2014, 16, 22882−22891. (13) Hansen, A. S.; Du, L.; Kjaergaard, H. G. Positively Charged Phosphorus as a Hydrogen Bond Acceptor. J. Phys. Chem. Lett. 2014, 5, 4225−4231. (14) Millen, D. J.; Zabicky, J. Hydrogen Bonding in Gaseous Mixtures. Part V. Infrared Spectra of Amine-Alcohol Systems. J. Chem. Soc. 1965, 0, 3080−3085. (15) Du, L.; Tang, S.; Hansen, A. S.; Frandsen, B. N.; Maroun, Z.; Kjaergaard, H. G. Subtle Differences in the Hydrogen Bonding of Alcohol to Divalent Oxygen and Sulfur. Chem. Phys. Lett. 2017, 667, 146−153. (16) Schrems, O.; Oberhoffer, H. M.; Luck, W. A. P. Hydrogen Bonding in Low-Temperature Matrixes: 1. Proton Donor Abilities of Fluoro Alcohols. Comparative Infrared Studies of ROH-O(CH3)2 Complex Formation in the Gas Phase, in Carbon Tetrachloride Solution, and in Solid Argon. J. Phys. Chem. 1984, 88, 4335−4342. (17) Biswal, H. S.; Shirhatti, P. R.; Wategaonkar, S. O-H···O versus O-H···S Hydrogen Bonding I: Experimental and Computational Studies on the p-Cresol···H2O and p-Cresol···H2S Complexes. J. Phys. Chem. A 2009, 113, 5633−5643. (18) Biswal, H. S.; Wategaonkar, S. O-H···O versus O-H···S Hydrogen Bonding. 3. IR-UV Double Resonance Study of Hydrogen Bonded Complexes of p-Cresol with Diethyl Ether and its Sulfur Analog. J. Phys. Chem. A 2010, 114, 5947−5957. (19) Biswal, H. S.; Shirhatti, P. R.; Wategaonkar, S. O-H···O versus O-H···S Hydrogen Bonding. 2. Alcohols and Thiols as Hydrogen Bond Acceptors. J. Phys. Chem. A 2010, 114, 6944−6955. (20) Bhaskar, K. R.; Bhat, S. N.; Singh, S.; Rao, C. N. R. Spectroscopic Studies of the Donor Properties of Triphenyl-Amine, -Phosphine, -Arsine and -Stibine Interaction with Iodine and Phenol. J. Inorg. Nucl. Chem. 1966, 28, 1915−1925. (21) Chojnowski, J. Triethylphosphine, Triethylarsine, and Triethylstibine as Hydrogen-Acceptors in Hydrogen Bonds. 2. Association

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b02523. Additional experimental spectra of the MeOH−pyridine, EtOH−pyridine and TFE−pyridine complexes; NCI isosurfaces for the MeOH, EtOH, and TFE complexes with TMA, DME, TMP, and DMS; selected optimized geometric parameters that describe the OH−Y intermolecular hydrogen bonds; comparison of NCI properties calculated for the gauche and trans conformers of the EtOH complexes (PDF)

■

Article

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +64-7-837-9391. ORCID

Joseph R. Lane: 0000-0002-4474-2941 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We thank K. H. Møller for providing helpful feedback on a draft version of the manuscript. We acknowledge financial support from the Danish Council for Independent Research− Natural Sciences. The Univ. of Waikato High Performance Computing Centre is gratefully acknowledged for computing time. 3459


Article

The Journal of Physical Chemistry A with Phenol and Methanol. B. Acad. Polym. Sci.-Chim. 1970, 18, 317− 327. (22) Rakova, L. V.; Romm, I. P.; Romanov, G. V.; Perepelkova, T. I.; Lapin, A. A.; Gol’dshtein, I. P.; Gur’yanova, E. N.; Pudovik, A. N. Formation Eenthalpies and Dipole Moments of the Complexes of Some Tertiary Phosphines with Trifluoroacetic Acid. Bull. Acad. Sci. USSR, Div. Chem. Sci. 1976, 25, 1315−1317. (23) Pogorelyi, V. K. Weak Hydrogen Bonds. Russ. Chem. Rev. 1977, 46, 316−336. (24) Épshtein, L. M.; Shubina, E. S.; Krylova, A. I.; Tolkunova, V. S.; Vil’chevskaya, V. D.; Kravtsov, D. N. Spectral Study of Ability of Ferrocenylphosphines and Ferrocenylphosphine Oxides to Form Hydrogen Bonds. Bull. Acad. Sci. USSR, Div. Chem. Sci. 1987, 36, 522−525. (25) Ryl’tsev, E. V.; Shurubura, A. K.; Egorov, Y. P.; Semenii, V. Y. Tertiary and Secondary Alkylphosphines as Partners in Hydrogen Bonding. Theor. Exp. Chem. 1981, 16, 384−390. (26) Broekmans, G.; Zeegers-Huyskens, T. Infrared Study (3700 3000 cm−1) of Hydrogen Bonds Involving Triphenyl Derivatives of Group V Elements. Spectrosc. Lett. 1982, 15, 515−525. (27) Sennikov, P. G. Weak Hydrogen-Bonding by Second-Row (PH3, H2S) and Third-Row (AsH3, H2Se) Hydrides. J. Phys. Chem. 1994, 98, 4973−4981. (28) Du, L.; Mackeprang, K.; Kjaergaard, H. G. Fundamental and Overtone Vibrational Spectroscopy, Enthalpy of Hydrogen Bond Formation and Equilibrium Constant Determination of the MethanolDimethylamine Complex. Phys. Chem. Chem. Phys. 2013, 15, 10194− 10206. (29) Johnson, E. R.; Keinan, S.; Mori-Sánchez, P.; Contreras-García, J.; Cohen, A. J.; Yang, W. Revealing Noncovalent Interactions. J. Am. Chem. Soc. 2010, 132, 6498−506. (30) Bader, R. Atoms in Molecules−A Quantum Theory; Oxford University Press: London, England, 1990. (31) Lane, J. R.; Contreras-García, J.; Piquemal, J. P.; Miller, B. J.; Kjaergaard, H. G. Are Bond Critical Points Really Critical for Hydrogen Bonding? J. Chem. Theory Comput. 2013, 9, 3263−3266. (32) Lane, J. R.; Schrøder, S. D.; Saunders, G. C.; Kjaergaard, H. G. Intramolecular Hydrogen Bonding in Substituted Aminoalcohols. J. Phys. Chem. A 2016, 120, 6371−6378. (33) Contreras-García, J.; Yang, W.; Johnson, E. R. Analysis of Hydrogen-Bond Interaction Potentials from the Electron Density: Integration of Noncovalent Interaction Regions. J. Phys. Chem. A 2011, 115, 12983−12990. (34) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A. et al. Gaussian 09; Gaussian, Inc.: Wallingford, CT, 2009. (35) Werner, H. J.; Knowles, P. J.; Knizia, G.; Manby, F. R.; Schütz, M. Molpro: A General-Purpose Quantum Chemistry Program Package. Wiley Interdiscip. Rev. Comput. Mol. Sci. 2012, 2, 242−253. (36) Zou, W. Molden2AIM. https://github.com/zorkzou/ Molden2AIM. (37) Contreras-García, J.; Johnson, E. R.; Keinan, S.; Chaudret, R.; Piquemal, J.-P.; Beratan, D. N.; Yang, W. NCIPLOT: A Program for Plotting Non-Covalent Interaction Regions. J. Chem. Theory Comput. 2011, 7, 625−632. (38) Saleh, G.; Lo Presti, L.; Gatti, C.; Ceresoli, D. NCImilano: An Electron-Density-Based Code for the Study of Noncovalent Interactions. J. Appl. Crystallogr. 2013, 46, 1513−1517. (39) Lu, T.; Chen, F. Multiwfn: A Multifunctional Wavefunction Analyzer. J. Comput. Chem. 2012, 33, 580−592. (40) Andersen, C. L.; Jensen, C. S.; Mackeprang, K.; Du, L.; Jorgensen, S.; Kjaergaard, H. G. Similar Strength of the NH−O and NH−S Hydrogen Bonds in Binary Complexes. J. Phys. Chem. A 2014, 118, 11074−11082. (41) Møller, K. H.; Hansen, A. S.; Kjaergaard, H. G. Gas Phase Detection of the NH-P Hydrogen Bond and Importance of Secondary Interactions. J. Phys. Chem. A 2015, 119, 10988−10998.

(42) Mukherjee, L. H.; Grunwald, E. Physical Properties and Hydrogen-bonding in the System Ethanol 2,2,2-Trifluoroethanol. J. Phys. Chem. 1958, 62, 1311−1314. (43) Hunter, E. P. L.; Lias, S. G. Evaluated Gas Phase Basicities and Proton Affinities of Molecules: an Update. J. Phys. Chem. Ref. Data 1998, 27, 413−656. (44) Heger, M.; Scharge, T.; Suhm, M. A.; et al. From Hydrogen Bond Donor to Acceptor: The Effect of Ethanol Fluorination on the First Solvating Water Molecule. Phys. Chem. Chem. Phys. 2013, 15, 16065. (45) Koch, U.; Popelier, P. L. A. Characterization of C-H-O Hydrogen Bonds on the Basis of the Charge Density. J. Phys. Chem. 1995, 99, 9747−9754. (46) Espinosa, E.; Souhassou, M.; Lachekar, H.; Lecomte, C. Topological Analysis of the Electron Density in Hydrogen Bonds. Acta Crystallogr., Sect. B: Struct. Sci. 1999, 55, 563−572. (47) Mata, I.; Alkorta, I.; Molins, E.; Espinosa, E. Universal Features of the Electron Density Distribution in Hydrogen-Bonding Regions: A Comprehensive Study Involving H−X (X = H, C, N, O, F, S, Cl, π) Interactions. Chem. - Eur. J. 2010, 16, 2442−2452. (48) Saleh, G.; Gatti, C.; Lo Presti, L. Energetics of Non-Covalent Interactions from Electron and Energy Density Distributions. Comput. Theor. Chem. 2015, 1053, 53−59. (49) Parrish, R. M.; Gonthier, J. F.; Corminbœuf, C.; Sherrill, C. D. Communication: Practical Intramolecular Symmetry Adapted Perturbation Theory via Hartree-Fock Embedding. J. Chem. Phys. 2015, 143, 051103. (50) Pastorczak, E.; Prlj, A.; Gonthier, J. F.; Corminbœuf, C. Intramolecular Symmetry-Adapted Perturbation Theory with a SingleDeterminant Wavefunction. J. Chem. Phys. 2015, 143, 224107. (51) Miller, B. J.; Lane, J. R.; Kjaergaard, H. G. Intramolecular OH···π interactions in alkenols and alkynols. Phys. Chem. Chem. Phys. 2011, 13, 14183−14193. (52) Lane, J. R.; Saunders, G. C.; Webb, S. J. Engineering of a Polar Crystal Structure Exclusively by π-π Stacking Between Aryl and Polyfluoroaryl Groups. CrystEngComm 2013, 15, 1293. (53) Thomsen, D. L.; Axson, J. L.; Schrøder, S. D.; Lane, J. R.; Vaida, V.; Kjaergaard, H. G. Intramolecular Interactions in 2-Aminoethanol and 3-Aminopropanol. J. Phys. Chem. A 2013, 117, 10260−10273. (54) Yan, Y.; Kuhn, O. Unraveling the Correlated Dynamics of the Double Hydrogen Bonds of Nucleic Acid Base Pairs in Solution. J. Phys. Chem. B 2011, 115, 5254−5259. (55) Pirc, G.; Mavri, J.; Novič, M.; Stare, J. Virtually Nonexistent Correlation between the OH Stretching Frequency and the Instantaneous Geometry in the Short Hydrogen Bond of Sodium Hydrogen Bis(sulfate): Advanced Chemometrics Analysis. J. Phys. Chem. B 2012, 116, 7221−7231.

3460


Kinetic Energy Density as a Predictor of Hydrogen-Bonded OH

Recommend Documents