DOI: 10.1002/cphc.201500324
Articles
Comparative Study of Halogen- and Hydrogen-Bond Interactions between Benzene Derivatives and Dimethyl Sulfoxide Yan-Zhen Zheng, Geng Deng, Yu Zhou, Hai-Yuan Sun, and Zhi-Wu Yu*[a] The halogen bond, similar to the hydrogen bond, is an important noncovalent interaction and plays important roles in diverse chemistry-related fields. Herein, bromine- and iodinebased halogen-bonding interactions between two benzene derivatives (C6F5Br and C6F5I) and dimethyl sulfoxide (DMSO) are investigated by using IR and NMR spectroscopy and ab initio calculations. The results are compared with those of interactions between C6F5Cl/C6F5H and DMSO. First, the interaction energy of the hydrogen bond is stronger than those of bromine- and chlorine-based halogen bonds, but weaker than iodine-based halogen bond. Second, attractive energies depend on 1/rn, in which n is between three and four for both
hydrogen and halogen bonds, whereas all repulsive energies are found to depend on 1/r8.5. Third, the directionality of halogen bonds is greater than that of the hydrogen bond. The bromine- and iodine-based halogen bonds are strict in this regard and the chlorine-based halogen bond only slightly deviates from 1808. The directional order is iodine-based halogen bond > bromine-based halogen bond > chlorine-based halogen bond > hydrogen bond. Fourth, upon the formation of hydrogen and halogen bonds, charge transfers from DMSO to the hydrogen- and halogen-bond donors. The CH3 group contributes positively to stabilization of the complexes.
1. Introduction Similar to the hydrogen atom, halogen atoms (X) in the form of R¢X can also be electron acceptors to interact with electron donors. These kinds of interactions are similar to hydrogen bonds (HBs) in some aspects and are called halogen bonds (XB).[1–5] Politzer’s group found that XB was a noncovalent interaction in which an electrophilic cap on a halogen atom, the s hole, attracted a nucleophilic site on an adjacent molecule.[6–10] A number of similarities have been found between XB and HB. For example, Metrangolo et al. found that XB also had directionality and selectivity.[11] Wang et al.[12] and Cheng et al.[13] found that there were also conventional redshifts of XB and unconventional blueshifts of XB. Lommerse et al.[14] and Umeyama and Morokuma[15] found that the formation of XB complexes was also closely related to electrostatic effects, polarization, and dispersion. Noh et al. found that both XB and HB could contribute positively to stabilization of the supramolecular structures.[16] Li et al. found that there was significant cooperativity between XB and HB in some complexes.[17] Nevertheless, the electronic structures of the halogen atoms are quite different from that of the hydrogen atom. There must be some differences in the two types of bonds.
[a] Y.-Z. Zheng, G. Deng, Y. Zhou, H.-Y. Sun, Prof. Z.-W. Yu Key Laboratory of Bioorganic Phosphorous Chemistry and Chemical Biology (Ministry of Education) Department of Chemistry, Tsinghua University Beijing 100084 (P.R. China) E-mail: [email protected] Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201500324.
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The interaction energies of XB and HB are different. The former can be greater or less than the latter.[15–17] In earlier work studying XB and HB between DMSO and three fluorobenzene derivatives (C6F5H, C6F5Cl, ClC6F4H), using both experimental and theoretical approaches, we found that HB was stronger than XB in these binary systems if chlorine was the halogen.[18] Considering that the strength of the bromine- (Br¢XB) and iodine-based halogen bonds (I¢XB) are normally greater than that of chlorine-based halogen bonds (Cl¢XB), we extend our work to these two cases herein. 1-Bromo-2,3,4,5,6-pentafluorobenzene (C6F5Br) and 1-iodo-2,3,4,5,6-pentafluorobenzene (C6F5I) were the selected XB donors, DMSO was still taken as the XB acceptor. Attenuated total reflectance Fourier transform infrared (ATR-FTIR) and 19F NMR spectroscopy and ab initio calculations were employed to investigate the halogen-bonding properties. In particular, excess IR spectroscopy[18–26] was used to reveal details of the molecular interactions. The directionality of XB is generally believed to be stronger than that of HB. For liquid systems, this understanding has been achieved mostly by quantum chemical calculations, in which the optimized XB complexes typically show near-linear configurations (the R¢X···Y angles are close to 1808), whereas the HB interaction pairs are more likely to be nonlinear (R¢ H···Y angles sometimes considerably less than 1808). The robustness of the collinearity of the intermolecular bond, however, has rarely been investigated. There have been a few publications on the angle dependences of the XB interaction energies.[27–31] Politzer’s group studied the interaction energy of three R¢Br···Y and R¢H···Y complexes as functions of the R¢H¢ Y and R¢Br¢Y angles.[27] They found that Br¢XB was more di-
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Articles rectional than HB. Recently, Hill and Legon studied HCCH···Y, ClH···Y, and FCl···Y complexes and found that the angle dependences were relatively flat for the hydrogen-bonding complexes, but there was significantly more pronounced curvature for halogen bonding.[28] Neither study involved XB when iodine was the halogen-bond donor. Clearly, more work is needed to address the issue. Herein, we have examined the angle dependences of the interaction energy of three XB pairs and one HB pair. We found that the energy increase slope was in the order of I¢XB > Br¢XB > Cl¢XB > HB, when the angle deviated from 1808. The distance dependence of the interaction energy of XB is another interesting academic question.[32–34] It is well known that, for normal dipole–dipole interactions in solution, the distance dependence of the attractive potential energy is 1/r6. In the case of hydrogen-bonding interactions, this is very complicated. The dependence varies from 1/r2 to 1/rx.[34] For XB, the only work, to the best of our knowledge, is from our recent publication. We found that the attractive XB energy between C6F5Cl and DMSO was (1/r3.1), whereas that between C6F5H and DMSO was (1/r3.3). Merz and Riley investigated the interaction energy of several halobenzene–formaldehyde complexes as a function of the halogen–oxygen separation distance.[32] It was found that the halogen-bond interaction energies increased as the size of the halogen atom increased. However, the distance-dependence equation was not provided. Herein, we examined the interaction energy between C6F5Br/C6F5I and DMSO and fitted the data to a function of Lennard–Jones form, E = a/rm + b/rn, and compared the results with those of the interaction pairs between C6F5Cl/C6F5H and DMSO. The dependence of the attractive term varied from 1/r3.1 to 1/r4.0, whereas that of the repulsive term was 1/r8.5 for all interaction pairs.
2. Results and Discussion 2.1. ATR and Excess IR Spectra Analysis First, we focus on the vs(C¢D) of [D6]DMSO, v(C¢Br) of C6F5Br, and v(C¢I) of C6F5I. The partial IR spectra of pure [D6]DMSO, DMSO, and the two benzene derivatives are presented in Figure 1. In Figure 1 A, the two bands centered at n˜ = 2124 and 2250 cm¢1 are attributed to the vs(C¢D) and vas(C¢D) of [D6]DMSO, respectively.[35] In Figure 1 B, the bands at n˜ 834 and 805 cm¢1 are attributed to v(C¢Br) in C6F5Br and v(C¢I) in C5F5I, respectively.[36] For [D6]DMSO, the band at n˜ 820 cm¢1 corresponds to the methyl rocking vibration.[35] To remove the influence of overlap, we used C6F5Br¢DMSO and C6F5I¢DMSO systems to analyze v(C¢Br) and v(C¢I). 2.1.1. C¢Br and C¢I Stretching Vibration in C6F5Br and C6F5I The IR and excess spectra in the v(C¢Br) and v(C¢I) regions of the C6F5Br¢DMSO and C6F5I¢DMSO systems are shown in Figures 2 and 3, respectively. A few features can be seen readily in the original IR spectra (Figures 2 A and 3 A) with increasing concentration of DMSO. First, v(C¢Br) and v(C¢I) gradually ChemPhysChem 2015, 16, 2594 – 2601
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Figure 1. Partial IR spectra of pure C6F5Br, C6F5I, DMSO, and [D6]DMSO in selected wavelength regions.
move to lower wavenumbers and the absorbance decreases. The wavenumber shifts of v(C¢Br) and v(C¢I) are shown in Figure S1 in the Supporting Information. The band positions of n(C¢Br) and v(C¢I) are redshifted by 1.1 and 3.5 cm¢1, respectively, which implies that the two single bonds are influenced by XBs and the halogen-bonding strength in the latter is greater. Second, the excess spectra in Figures 2 B and 3 B contain positive and negative bands and the positions of the positive bands are nearly fixed. The negative bands are at higher wavenumbers, whereas the positive bands are at lower wavenumbers compared with v(C¢Br) and v(C¢I) in pure C6F6Br and C6F5I, respectively. In the C6F5Br¢DMSO system, if the molar fraction of DMSO is less than 0.5, there are two (or even more) negative bands and one positive band. If the molar fraction of DMSO is greater than 0.5, the negative band is apparently singular (Figure 2 B). In the C6F5I¢DMSO system, if the molar fraction of DMSO is less than 0.3, there are also two negative bands and one positive band. If the molar fraction of DMSO is greater than 0.3, the negative bands look like singular ones (Figure 3 B). In an excess spectrum, the negative and positive bands represent the possible disappearing/increasing and appearing/increasing forms of the relevant species in the mixture compared with the pure liquid. In the lower concentration region of DMSO, there are two negative bands in v(C¢Br) and
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Figure 2. ATR-FTIR (A) and excess IR spectra (B) of the C6F5Br¢DMSO system in the range of v(C¢Br). From top to bottom: the molar fraction of DMSO increases from 0 to 1 in A) with an increment of about 0.1. The dashed and dash–dotted lines depict spectra of pure C6F5Br and DMSO in A). The inset model in A) is a sketch of C6F5Br self-association.
v(C¢I), which indicates that the amounts of the two species in the mixture decrease compared with those in the pure liquid. At higher concentrations of DMSO, the phenomenon is less clear. However, it is noted that the negative band is located between the two negative bands at lower concentrations of DMSO, which indicates that the broad, negative band at higher concentrations of DMSO is not a single band and may be composed of the two former negative bands. The two negative bands in the excess spectra indicate that there are at least two species in solutions of pure C6F6Br and C6F5I. The two species may be the C6F5Br/C6F5I monomer and C6F5Br/C6F5I selfassociated complex. Take C6F5Br self-association for example: the combination of C6F5Br monomers is through p–p stacking. A sketch of C6F5Br self-association is shown as the inset in Figure 2 A. The wavenumbers of v(C¢Br) and v(C¢I) in C6F5Br and C5F5I self-associates are lower than those of the C6F5Br and C6F5I monomers. When either of the solutions was diluted with DMSO, the amount of self-associates decreased to form C6F5Br¢DMSO and C6F5I¢DMSO complexes. The two negative bands in the excess spectra indicate the decrease in C6F5Br/ C6F5I monomer and C6F5Br/C6F5I self-associate, whereas the positive band over the entire concentration range indicates ChemPhysChem 2015, 16, 2594 – 2601
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Figure 3. ATR-FTIR (A) and excess IR spectra (B) of the C6F5I¢DMSO system in the range of v(C¢I). From top to bottom: the molar fraction of DMSO increases from 0 to 1 in A) with an increment of about 0.1. The dashed and dash–dotted lines depict spectra of pure C6F5I and DMSO in A).
the appearance of C6F5Br/C6F5I¢DMSO complexes. The wavenumber shifts of appearing species are lower than those of disappearing species. These results have been confirmed by quantum chemical calculations (see Section 2.3). An interesting feature in Figures 2 B and 3 B is the relative intensity of the negative bands when the molar fraction of DMSO is about 0.1. As indicated by the arrows in Figure 2 B, the intensity of the negative band in the higher wavenumber region is stronger than that in the lower wavenumber region. In Figure 3 B, however, the relative intensities of the two negative bands are opposite. This may be attributed to stronger interactions between C6F5I and DMSO than those between C6F5Br and DMSO. Thus, it is easier for DMSO to break apart the C6F5I self-associate, resulting in the stronger negative intensity of the band representing the self-associate (Figure 3 B). For the predominant breakup of the self-associate of C6F5Br, it requires a higher fraction of DMSO (0.2 and above). 2.1.2. vs(C¢D) in [D6]DMSO Shown in Figure 4 are the ATR-FTIR and respective excess spectra of vs(C¢D) in the C6F5Br¢[D6]DMSO and C6F5I¢[D6]DMSO systems over the entire concentration range. Two clear features
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Articles ple, at this concentration, the area of the negative band in the excess spectra of the C6F5Br¢[D6]DMSO system is barely seen, whereas that of C6F5I¢[D6]DMSO is quite pronounced. In combination with our former work,[18] for the C6F5H¢[D6]DMSO and ClC6F4H¢[D6]DMSO systems, the relative intensities of the negative bands are weaker than those of the C6F5I¢[D6]DMSO system. For the C6F5Cl¢[D6]DMSO system, the negative band does not appear until a molar fraction, x(C6F5Cl), of about 0.4. This indicates that C6F5I is a more powerful benzene derivative to break apart the DMSO self-associates, and the possible sequential order of the interaction strength of the five systems is as follows: C6F5I¢DMSO > C6F5H¢DMSO ClC6F4H¢DMSO > C6F5Br¢DMSO > C6F5Cl¢DMSO. This conclusion is supported by the quantum chemical calculation reported in Section 2.3.1. 2.2.
19
F NMR Spectroscopy Analysis
19
Figure 4. ATR-FTIR (A, B) and excess IR (C, D) spectra of the C6F5Br¢ [D6]DMSO (A, C) and C6F5I¢[D6]DMSO (B, D) systems in the range of vs(C¢D). From top to bottom: the molar fraction of [D6]DMSO decreases from 1 to 0 in A) and B) with a decrement of about 0.1. The dashed and dash–dotted lines depict spectra of pure [D6]DMSO, C6F5Br, and C6F5I.
can be seen readily in the original IR spectra (Figure 4 A and B): the absorbance of vs(C¢D) decreases monotonically and the bands gradually shift to higher wavenumber with increasing benzene derivatives. The wavenumber shifts of vs(C¢D) are shown in Figure S2 in the Supporting Information. Excess IR spectra provide another angle to observe molecular interactions between DMSO and the halogen-containing benzene derivatives, and the results are shown in Figure 4 C and D. General features of the two excess spectra are positive bands at higher wavenumbers and negative bands at lower wavenumbers in the two systems. With increasing molar fraction of benzene derivatives, the ratio of the absolute area of the negative band to the positive band increases gradually. The positions of positive and negative bands are nearly fixed. This means that DMSO molecules may have formed fixed interaction complexes in the solution. They are attributed to the absorptions of the [D6]DMSO molecules complexing with C6F5Br/C6F5I and those in self-associates. Another feature in Figure 4 C and D is that the areas of the negative bands in the excess spectra are different for the two systems. Taking the molar fraction of x(C6F5X) 0.1 as an examChemPhysChem 2015, 16, 2594 – 2601
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F NMR spectroscopy measurements of pure C6F5Br and C6F5I, and a series of C6F5Br¢DMSO, and C6F5I¢DMSO mixtures, were performed at 298 K, and the changes to the chemical shift of individual fluorine atoms during the dilution process were evaluated. Assignments of the 19F NMR spectroscopy signals to the fluorine atoms in C6F5Br and C6F5I are displayed in Figure 5. The concentration dependences of the spectra are displayed in Figure 6 A and B, and the dependences of chemical shift variations (Dd) of the two benzene derivatives on the molar fraction of DMSO are displayed in Figure S3 in Supporting Information. As observed, all chemical shift values become more and more negative with increasing concentration of DMSO, which means that the fluorine atoms gradually move upfield during the dilution process. This can be explained by the formation of halogen-bonding complexes because it is known that, upon forming XBs, electrons transfer from acceptors to donors, which causes upfield movement of the chemical shifts of the groups in the donors.[37, 38] Detailed charge gains of the fluorine atoms in the two benzene derivatives upon forming XB with DMSO have been calculated by the quantum chemical method, by using a 1:1 ratio in vacuum. The results are marked on the molecules in Figure 5. The negative values indi-
Figure 5. 19F NMR spectroscopy assignments of neat C6F5Br and C6F5I.
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Articles directional than XB. This is in agreement with conclusions reported in the literature.[40] Table 1 lists selected geometrical and vibrational parameters of the investigated complexes. For a better comparison, the corresponding data of the C6F5H¢DMSO and C6F5Cl¢DMSO systems is also listed.[18] As can be seen from these results, the C¢ Br and C¢I bonds in C6F5Br and C6F5I, respectively, are elongated upon complexation. The elongations are concomitant with a decrease in the stretching frequencies (redshift) and an increase in the IR intensities of the respective bonds. These are in agreement with the experimental results discussed in previous sections.
Table 1. Selected bond lengths (r in æ), vibrational frequencies (n in cm¢1), and IR intensities (I in km mol¢1) of the optimized complexes C6F5X¢DMSO (X = H, Cl, Br, or I). Data in parentheses represent changes in bond length (r), vibrational frequency (n), and intensity (I) upon complexation.
Complex
r(C¢X)
v(C¢X)
I(C¢X)
C6F5I¢DMSO
2.096 (0.016) 1.877 (0.004) 1.711 (0.002) 1.091 (0.007)
809.99 (¢1.81) 839.24 (¢1.71) 887.43 (¢0.14) 3157.67 (¢98.47)
72.83 (12.90) 109.79 (10.75) 151.14 (4.74) 318.13 (310.9)
C6F5Br¢DMSO C6F5Cl¢DMSO18
Figure 6. The molar fraction dependences of the 19F NMR spectroscopy signals in the C6F5Br¢DMSO (A) and C6F5I¢DMSO (B) systems.
C6F5H¢DMSO18
cate that all fluorine atoms in the two benzene derivatives gain charge upon forming XBs with DMSO.
2.3. Quantum Chemical Calculations 2.3.1. Structures The optimized geometries of individual molecules of DMSO, C6F5Br, and C6F5I, as well as the C6F5Br¢DMSO and C6F5I¢DMSO complexes, were studied. The sum of van der Waals atomic radii of bromine and oxygen (3.35 æ) and that of iodine and oxygen (3.55 æ) were used as critical values to determine the presence of a halogen bond.[39] The inset in Figure 7 shows the optimized geometries of the two complexes. The corresponding interaction energies of the complexes are displayed below the structures. The O¢X separations demonstrate that XBs form in both cases. Based on the interaction energies shown in Figure 7 and previous work,[18] we can easily obtain the sequential order of the interaction strength: C6F5I¢DMSO > C6F5H¢DMSO > C6F5Br¢DMSO > C6F5Cl¢DMSO. The strength of HB is between those of I¢XB and Br¢XB. Regarding the directionality of the halogen bonds, as shown in Figure 7, the optimized C¢Br···O and C¢I···O angles are all 1808. For C¢Cl···O angles,[18] they are 179 and 1768, which are very close to 1808. This reflects the characteristic directionality of XBs. In comparison, the optimized C¢H···O angles, as in the complexes of ClC6F4H¢DMSO and C6F5H¢DMSO, are 167 and 1638, which are much less than 1808 and show that HB are less ChemPhysChem 2015, 16, 2594 – 2601
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2.3.2. Angle Dependence of Energy To gain a deeper understanding of the directionality of halogen and hydrogen bonds, the dependences of the energies of XB and HB on the angle C¢X¢O (X = I, Br, Cl, and H) were studied. Single-point energy changes of the four complexes (C6F5H¢DMSO, C6F5Cl¢DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO)
Figure 7. Distance dependences of the interaction energies of the C6F5Cl¢ DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO complexes through halogen-bonding interactions. The inset shows the optimized geometries and corresponding interaction energies of the complexes. XB bonds are denoted by dashed lines, and the corresponding XB distances [æ] and bond angles [8] are labeled.
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Articles interactive nature changes from repulsive to attractive. There is a position with the lowest energy in each case that corresponds to the most stable optimized geometry shown in the inset of Figure 8. The data can be fitted to a function of Lennard–Jones form: E = a/rm¢b/rn. For Br¢XB, m = 8.5 and n = 3.2, and for I¢XB, m = 8.5 and n = 4.0. Combined with previous work,[18] the attractive energy of HB decays slightly slower than that of the I¢XB, although it is quicker than those of Br¢XB and Cl¢XB. All repulsive energies of HB and XB decay at 1/r8.5. 2.3.4. NBO Analysis Charge distributions of the most stable halogen-bonded complexes and the respective single molecules were examined by means of natural bond orbital (NBO) analysis. Some of the results are listed in Table 2. To compare HB and XB more clearly, the corresponding results for C6F5H¢DMSO and C6F5Cl¢DMSO are also listed in Table 2.[18]
Figure 8. XB and HB angle dependences of the interaction energies of the complexes C6F5H¢DMSO, C6F5Cl¢DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO.
in their most stable forms as functions of the bond angle were calculated, and the results are shown in Figure 8 with the corresponding angle shown in the inset. The optimized positions are taken as the reference point (08 in Figure 8). The angle change is expressed by the angle differences between the most stable optimized structure and the analytical structures. Thus, with decreasing angle from the optimized position, the energies increase monotonically, which indicates that the interaction strength gradually reduces. More importantly, the angle dependences of XB or HB reflect the robustness of directionality. Thus, a decrease of 208 in the XB angle of C6F5I¢DMSO needs an energy compensation of about 6.7 kJ mol¢1. For the HB in C6F5H¢DMSO, however, the energy compensation is only about 1.6 kJ mol¢1. It is clear, therefore, that the directionality of the XB angles in C6F5X¢DMSO are very robust (ca. 1808). On the other hand, the directionality of the HB in C6F5H¢DMSO can be perturbed easily; a deviation of 138 from the straightline arrangement of the bond angle in C6F5H¢DMSO is observed.[18] Thus, the curves in Figure 8 demonstrate that the directionality of the three XBs follow the sequential order of I¢XB > Br¢ XB > Cl¢XB. In addition, all XBs have better directionality than that of HB. This can be explained by the fact that the positive electrostatic potential of a halogen atom is on the outermost portion of the atom, centered around the intersection of its surface with the C¢X axis (s hole), whereas the positive electrostatic potential of a hydrogen atom covers a much larger area of the atomic surface.[18] 2.3.3. Distance Dependence of Energy To study the dependence of XB energies on bond length, single-point energies of the two complexes in their most stable forms as functions of the Br···O and I···O separations were calculated, and the results are shown in Figure 7. The corresponding result for C6F5Cl¢DMSO is also drawn in the Figure 7.[18] As observed, with increasing separation, the energies change from positive to negative, which indicates that the ChemPhysChem 2015, 16, 2594 – 2601
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Table 2. Charges (q in e) on selected atoms and groups and charge change (CC) of the entire DMSO molecule. Data in parentheses represent CC values of the corresponding atoms or groups upon complexion. Negative values indicate increases in charge density. Complex
Methyl
X
O
CC
C6F5I¢DMSO
¢0.199 (0.038) ¢0.211 (0.026) ¢0.215 (0.022) ¢0.202 (0.035)
0.345 (0.055) 0.241 (0.043) 0.133 (0.034) 0.288 (0.046)
¢0.985 (¢0.031) ¢0.970 (¢0.016) ¢0.966 (¢0.012) ¢0.979 (¢0.025)
0.021
C6F5Br¢DMSO C6F5Cl¢DMSO18 C6F5H¢DMSO18
0.011 0.009 0.018
As observed from the results in Table 2, halogen-bonding donor molecules gain charges (negative charge) upon complexation, with markedly larger gain in the case of C6F5I¢ DMSO. Combined with former work,[18] the charge gains follow the sequential order of C6F5I¢DMSO > C6F5H¢DMSO > C6F5Br¢ DMSO > C6F5Cl¢DMSO. Clearly, stronger interactions result in the XB or HB donors having more charge. The halogen atoms, however, show a decrease in electron density, similar to a hydrogen atom when it acts as the HB donor. The oxygen atoms in the five complexes all gain charge upon complexation. In addition, the methyl groups of DMSO in all five complexes lose negative charge upon complexation. Stronger interactions result in greater charge loss from the methyl groups. Altogether, there is charge transfer from the hydrogen/halogen-bond acceptors (DMSO) to the donors. Although the methyl groups of DMSO do not participate directly in the formation of the halogen/hydrogen bonds, they donate electrons to the S=O group and contribute positively to the stability of the complexes. This is in agreement with our earlier work on the back groups.[18–20]
3. Conclusions By employing quantum chemical calculation and spectroscopic techniques, and taking several halogen-substituted benzene
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Articles derivatives, we performed a comparative study on XBs and HBs. Several conclusions could be drawn: 1) XBs were more directional than HBs. The optimized conformations of the interaction complexes in vacuum showed that Br¢XB and I¢XB were strict in this regard and Cl¢XB only slightly deviated from 1808. The angle dependences of the interaction energies further demonstrated that the directional order was I¢XB > Br¢ XB > Cl¢XB > HB. 2) Upon complexation, the C¢H and C¢X groups in benzene derivatives were all elongated, concomitant with decreases in the corresponding stretching frequencies (redshift) and increases in the IR intensities. 3) The interaction energy of the three XBs decreased with decreasing mass of the halogen atoms in the benzene derivatives. The interaction energy of HB was between those of Br¢XB and I¢XB. 4) The separation distance dependence of the interaction energies were described in the form of a Lennard–Jones equation. The attractive energies depended on 1/r3.3 for HB, and 1/r4.0, 1/r3.2, and 1/r3.1 for I¢XB, Br¢XB, and Cl¢XB, respectively; all repulsive energies of XB and HB decayed with 1/r8.5. Thus, stronger attractive interactions resulted in faster energy decays. This was true for HB and the three XBs. 5) Upon the formation of HB or XB, electrons transferred from DMSO to the benzene derivatives. Although the methyl groups of DMSO did not participate directly in the formation of HB/XB, they donated negative charge to the S=O group and contributed positively to the stability of the complexes and were not inert. Comparing the four interaction pairs, we found that stronger interactions made the methyl group lose more charge. These comparison studies on the interactions between fluorobenzene derivatives and DMSO provided in-depth information to understand the differences between XB and HB interactions in solution.
Experimental Section Materials C6F5Br (> 98 %) and C6F5I (> 98 %) were purchased from J&K. [D6]DMSO (> 99.8 % of deuterium) was purchased from Cambridge Isotopes Laboratories (USA). DMSO (> 99.5 %) was obtained from Beijing Chemical Plant (Beijing, P.R. China).
FTIR Spectroscopy FTIR spectra over the range from n˜ = 4000 to 650 cm¢1 were collected at room temperature ( 25 8C) by using a Nicolet 5700 FTIR spectrometer, equipped with a deuterated triglycine sulfate (DTGS) detector. The ATR cells, made of trapezoidal ZnSe and Ge crystals with incident angles of 45 and 608, corresponding to 12 and 7 reflections, respectively, were employed. The Ge crystal, with fewer numbers of reflections, and thus, shorter effective light path, was used to examine the strong stretching bands of C¢Br and C¢I. The ZnSe crystal was used to examine the weak stretching bands of C¢ D. All spectra were recorded with a resolution of 2 cm¢1, 16 parallel scans, and a zero filling factor of 2. For each sample, three parallel measurements were performed. The refractive indices of solutions were measured with an Abbe refractometer at 25 8C. The formula suggested by Hansen[41] was used to do the ATR corrections. ChemPhysChem 2015, 16, 2594 – 2601
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Excess IR Spectroscopy The theory of excess IR spectroscopy has been described in detail elsewhere.[19, 20] Briefly, an excess IR spectrum is defined as the difference between the spectrum of a real solution and that of the respective ideal solution under identical conditions. The working equation for calculating the excess IR spectrum is given by Equation (1): eE ¼
¨ ¦ A ¢ x1 e*1 þ x2 e*2 dðC1 þ C2 Þ
ð1Þ
in which A is the absorbance of the mixture; d is the light path length; Ci, xi, and e*i are the molarity, molar fraction, and molar absorption coefficient, respectively, of component i in its pure state. Matlab 7.0 (Math Works Inc., Natick, MA) was used to manipulate spectral data pretreatment, namely, the subtraction, truncation, and baseline correction, to calculate excess IR spectra and to integrate the bands in the excess IR spectra.
NMR Spectroscopy Measurements 19
F NMR spectroscopy is a sensitive indicator to provide intermolecular interaction information involving fluorine-substituted molecules, owing to its large chemical shift range and high intrinsic sensitivity (85 % of 1H NMR spectroscopy). The 19F NMR spectroscopy measurements were performed on a JEOL JNM-ECA 600 NMR (600 MHz) spectrometer at 298 K. Pure CF3COOH was used as an external standard to avoid its influence on the chemicals in the binary mixtures.
Quantum Chemical Calculations All computations were performed with the Gaussian 03 program.[42] We used the second-order many-body perturbation theory (MP2) with mixed basis sets to perform quantum chemical computations. The molecular energy, geometry, vibrational frequency, vibrational intensity, and natural population analysis (NPA) charge of all isolated single molecules and their complexes were fully optimized with the LANL2DZ[43] basis set for bromine and iodine atoms, and the aug-cc-pVDZ[44] basis set for hydrogen, carbon, fluorine, sulfur, and oxygen atoms. This allowed consideration of the relativistic effects in bromine and iodine atoms. The optimized geometries at the local energy minimum were ensured by the absence of imaginary vibrational frequency. The interaction energy was estimated as the difference between the total energy of the complex and the sum of total energies of the two individual molecules. Meanwhile, the basis set superposition error (BSSE) correction computed through the counterpoise method of Boys and Bernardi was estimated to obtain accurate interaction energies of the complexes.[45] The NPA charges were obtained by using the NBO approach.[46] To evaluate the distance dependence of the energy of XB, single-point energies were calculated by changing the distance between an electron donor (oxygen atom in DMSO) and an electron acceptor (bromine atom in C6F5Br and iodine atom in C6F5I).
Acknowledgements This work was supported by the Natural Science Foundation of China (grant nos. 21473099, 21273130).
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Articles Keywords: halogen bonds · hydrogen bonds · IR spectroscopy · structure–activity relationships · substituent effects [1] R. H. Crabtree, P. E. M. Siegbahn, O. Eisenstein, A. L. Rheingold, T. F. Koetzle, Acc. Chem. Res. 1996, 29, 348 – 354. [2] A. C. Legon, Phys. Chem. Chem. Phys. 2010, 12, 7736 – 7747. [3] A. Priimagi, G. Cavallo, P. Metrangolo, G. Resnati, Acc. Chem. Res. 2013, 46, 2686 – 2695. [4] M. Erd¦lyi, Chem. Soc. Rev. 2012, 41, 3547 – 3557. [5] A. C. Legon, Angew. Chem. Int. Ed. 1999, 38, 2686 – 2714; Angew. Chem. 1999, 111, 2850 – 2880. [6] T. Clark, M. Hennemann, J. S. Murray, P. Politzer, J. Mol. Model. 2007, 13, 291 – 296. [7] P. Politzer, P. Lane, M. C. Concha, Y. Ma, J. S. Murray, J. Mol. Model. 2007, 13, 305 – 311. [8] P. Politzer, J. S. Murraya, T. Clark, Phys. Chem. Chem. Phys. 2010, 12, 7748 – 7757. [9] P. Politzer, J. S. Murraya, ChemPhysChem 2013, 14, 278 – 294. [10] G. Trogdon, J. S. Murray, M. C. Concha, P. Politzer, J. Mol. Model. 2007, 13, 313 – 318. [11] P. Metrangolo, H. Neukirch, T. Pilati, G. Resnati, Acc. Chem. Res. 2005, 38, 386 – 395. [12] W. Z. Wang, N. B. Wong, W. X. Zheng, A. M. Tian, J. Phys. Chem. A 2004, 108, 1799 – 1805. [13] J. B. Cheng, R. Li, Q. Z. Li, B. Jing, Z. B. Liu, W. Z. Li, B. A. Gong, J. Z. Sun, J. Phys. Chem. A 2010, 114, 10320 – 10325. [14] J. P. M. Lommerse, A. J. Stone, R. Taylor, F. H. Allen, J. Am. Chem. Soc. 1996, 118, 3108 – 3116. [15] H. Umeyama, K. Morokuma, J. Am. Chem. Soc. 1977, 99, 1316 – 1332. [16] S. K. Noh, J. H. Jeon, W. J. Jang, H. Kim, S. H. Lee, M. W. Lee, J. Lee, S. Han, S. Kahng, ChemPhysChem 2013, 14, 1177 – 1181. [17] Q. Z. Li, Q. Q. Lin, W. Z. Li, J. B. Cheng, B. A. Gong, J. Z. Sun, ChemPhysChem 2008, 9, 2265 – 2269. [18] Y. Z. Zheng, N. N. Wang, Y. Zhou, Z. W. Yu, Phys. Chem. Chem. Phys. 2014, 16, 6946 – 6956. [19] Q. Z. Li, G. S. Wu, Z. W. Yu, J. Am. Chem. Soc. 2006, 128, 1438 – 1439. [20] Q. Z. Li, N. N. Wang, Q. Zhou, S. Q. Sun, Z. W. Yu, Appl. Spectrosc. 2008, 62, 166 – 170. [21] Y. Koga, F. Sebe, T. Minami, K. Otake, K. Saitow, K. Nishikawa, J. Phys. Chem. B 2009, 113, 11928 – 11935. [22] Y. Feng, D. Rainteau, C. Chachaty, Z. W. Yu, C. Wolf, P. J. Quinn, Biophys. J. 2004, 86, 2208 – 2217. [23] Y. Z. Zheng, N. N. Wang, J. J. Luo, Y. Zhou, Z. W. Yu, Phys. Chem. Chem. Phys. 2013, 15, 18055 – 18064. [24] N. N. Wang, Y. Wang, H. F. Cheng, Z. Tao, J. Wang, W. Z. Wu, RSC Adv. 2013, 3, 20237 – 20245.
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[25] J. Kiefer, M. M. Molina, K. Noack, ChemPhysChem 2012, 13, 1213 – 1220. [26] S. Corsetti, F. M. Zehentbauer, D. McGloin, J. Kiefer, Fuel 2015, 141, 136 – 142. [27] Z. P. Shields, J. S. Murray, P. Politzer, Int. J. Quantum Chem. 2010, 110, 2823 – 2832. [28] J. G. Hill, A. C. Legon, Phys. Chem. Chem. Phys. 2015, 17, 858 – 867. [29] R. Wilcken, M. O. Zimmermann, A. Lange, S. Zahn, B. Kirchner, F. M. Boeckler, J. Chem. Theory Comput. 2011, 7, 2307 – 2315. [30] D. Dolenc, B. Modec, New J. Chem. 2009, 33, 2344 – 2349. [31] H. S. El-Sheshtawy, B. S. Bassil, K. I. Assaf, U. Kortz, W. M. Nau, J. Am. Chem. Soc. 2012, 134, 19935 – 19941. [32] K. E. Riley, K. M. Merz, J. Phys. Chem. A 2007, 111, 1688 – 1694. [33] R. Li, Q. Z Li, J. B. Cheng, Z. B. Liu, W. Z. Li, ChemPhysChem 2011, 12, 2289 – 2295. [34] J. N. Israelachvili, Intermolecular and Surface Forces, 3rd ed., Academic Press, San Diego, 2011. [35] W. D. Horrocks, Jr., F. A. Cotton, Spectrochim. Acta 1961, 17, 134 – 147. [36] S. G. Frankiss, D. J. Harrison, Spectrochim. Acta 1975, 31, 1839 – 1864. [37] D. Hauchecorne, B. J. van der Veken, W. A. Herrebout, P. E. Hansen, Chem. Phys. 2011, 381, 5 – 10. [38] P. Metrangolo, W. Panzeri, F. Recupero, G. Resnati, J. Fluorine Chem. 2002, 114, 27 – 33. [39] L. Pauling, The Nature of the Chemical Bond, 3rd ed., Cornell University Press, New York, 1960. [40] C. Wang, D. Danovich, Y. Mo, S. Shaik, J. Chem. Theory Comput. 2014, 10, 3726 – 3737. [41] W. N. Hansen, Spectrochim. Acta 1965, 21, 815 – 833. [42] Gaussian 03, Revision B.03, M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A. Robb, J. R. Cheeseman, V. G. Zakrzewski, J. A. Montgomery, Jr., R. E. Stratmann, J. C. Burant, S. Dapprich, J. M. Millam, A. D. Daniels, K. N. Kudin, M. C. Strain, O. Farkas, J. Tomasi, V. Barone, M. Cossi, R. Cammi, B. Mennucci, C. Pomelli, C. Adamo, S. Clifford, J. Ochterski, A. Petersson, P. Y. G. Ayala, Q. Cui, K. Morokuma, D. K. Malick, A. D. Rabuck, K. Raghavachari, J. B. Foresman, J. Cioslowski, J. V. Ortiz, A. G. Baboul, B. B. Stefanov, G. Liu, A. Liashenko, P. Piskorz, I. Komaromi, R. Gomperts, R. L. Martin, D. J. Fox, T. Keith, M. A. Al-Laham, C. Y. Peng, A. Nanayakkara, C. Gonzalez, M. Challacombe, P. M. W. Gill, B. Johnson, W. Chen, M. W. Wong, J. L. Andres, C. Gonzalez, M. Head-Gordon, E. S. Replogle, J. A. Pople, Gaussian Inc., Pittsburgh, PA, 2003. [43] W. R. Wadt, P. J. Hay, J. Chem. Phys. 1985, 82, 270 – 283. [44] T. H. Dunning, J. Chem. Phys. 1989, 90, 1007 – 1023. [45] S. F. Boys, F. Bernardi, Mol. Phys. 1970, 19, 553 – 566. [46] A. E. Reed, L. A. Curtiss, F. Weinhold, Chem. Rev. 1988, 88, 899 – 926.
Received: April 18, 2015 Published online on June 26, 2015
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Comparative Study of Halogen- and Hydrogen-Bond Interactions between Benzene Derivatives and Dimethyl Sulfoxide Yan-Zhen Zheng, Geng Deng, Yu Zhou, Hai-Yuan Sun, and Zhi-Wu Yu*[a] The halogen bond, similar to the hydrogen bond, is an important noncovalent interaction and plays important roles in diverse chemistry-related fields. Herein, bromine- and iodinebased halogen-bonding interactions between two benzene derivatives (C6F5Br and C6F5I) and dimethyl sulfoxide (DMSO) are investigated by using IR and NMR spectroscopy and ab initio calculations. The results are compared with those of interactions between C6F5Cl/C6F5H and DMSO. First, the interaction energy of the hydrogen bond is stronger than those of bromine- and chlorine-based halogen bonds, but weaker than iodine-based halogen bond. Second, attractive energies depend on 1/rn, in which n is between three and four for both
hydrogen and halogen bonds, whereas all repulsive energies are found to depend on 1/r8.5. Third, the directionality of halogen bonds is greater than that of the hydrogen bond. The bromine- and iodine-based halogen bonds are strict in this regard and the chlorine-based halogen bond only slightly deviates from 1808. The directional order is iodine-based halogen bond > bromine-based halogen bond > chlorine-based halogen bond > hydrogen bond. Fourth, upon the formation of hydrogen and halogen bonds, charge transfers from DMSO to the hydrogen- and halogen-bond donors. The CH3 group contributes positively to stabilization of the complexes.
1. Introduction Similar to the hydrogen atom, halogen atoms (X) in the form of R¢X can also be electron acceptors to interact with electron donors. These kinds of interactions are similar to hydrogen bonds (HBs) in some aspects and are called halogen bonds (XB).[1–5] Politzer’s group found that XB was a noncovalent interaction in which an electrophilic cap on a halogen atom, the s hole, attracted a nucleophilic site on an adjacent molecule.[6–10] A number of similarities have been found between XB and HB. For example, Metrangolo et al. found that XB also had directionality and selectivity.[11] Wang et al.[12] and Cheng et al.[13] found that there were also conventional redshifts of XB and unconventional blueshifts of XB. Lommerse et al.[14] and Umeyama and Morokuma[15] found that the formation of XB complexes was also closely related to electrostatic effects, polarization, and dispersion. Noh et al. found that both XB and HB could contribute positively to stabilization of the supramolecular structures.[16] Li et al. found that there was significant cooperativity between XB and HB in some complexes.[17] Nevertheless, the electronic structures of the halogen atoms are quite different from that of the hydrogen atom. There must be some differences in the two types of bonds.
[a] Y.-Z. Zheng, G. Deng, Y. Zhou, H.-Y. Sun, Prof. Z.-W. Yu Key Laboratory of Bioorganic Phosphorous Chemistry and Chemical Biology (Ministry of Education) Department of Chemistry, Tsinghua University Beijing 100084 (P.R. China) E-mail: [email protected] Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201500324.
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The interaction energies of XB and HB are different. The former can be greater or less than the latter.[15–17] In earlier work studying XB and HB between DMSO and three fluorobenzene derivatives (C6F5H, C6F5Cl, ClC6F4H), using both experimental and theoretical approaches, we found that HB was stronger than XB in these binary systems if chlorine was the halogen.[18] Considering that the strength of the bromine- (Br¢XB) and iodine-based halogen bonds (I¢XB) are normally greater than that of chlorine-based halogen bonds (Cl¢XB), we extend our work to these two cases herein. 1-Bromo-2,3,4,5,6-pentafluorobenzene (C6F5Br) and 1-iodo-2,3,4,5,6-pentafluorobenzene (C6F5I) were the selected XB donors, DMSO was still taken as the XB acceptor. Attenuated total reflectance Fourier transform infrared (ATR-FTIR) and 19F NMR spectroscopy and ab initio calculations were employed to investigate the halogen-bonding properties. In particular, excess IR spectroscopy[18–26] was used to reveal details of the molecular interactions. The directionality of XB is generally believed to be stronger than that of HB. For liquid systems, this understanding has been achieved mostly by quantum chemical calculations, in which the optimized XB complexes typically show near-linear configurations (the R¢X···Y angles are close to 1808), whereas the HB interaction pairs are more likely to be nonlinear (R¢ H···Y angles sometimes considerably less than 1808). The robustness of the collinearity of the intermolecular bond, however, has rarely been investigated. There have been a few publications on the angle dependences of the XB interaction energies.[27–31] Politzer’s group studied the interaction energy of three R¢Br···Y and R¢H···Y complexes as functions of the R¢H¢ Y and R¢Br¢Y angles.[27] They found that Br¢XB was more di-
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Articles rectional than HB. Recently, Hill and Legon studied HCCH···Y, ClH···Y, and FCl···Y complexes and found that the angle dependences were relatively flat for the hydrogen-bonding complexes, but there was significantly more pronounced curvature for halogen bonding.[28] Neither study involved XB when iodine was the halogen-bond donor. Clearly, more work is needed to address the issue. Herein, we have examined the angle dependences of the interaction energy of three XB pairs and one HB pair. We found that the energy increase slope was in the order of I¢XB > Br¢XB > Cl¢XB > HB, when the angle deviated from 1808. The distance dependence of the interaction energy of XB is another interesting academic question.[32–34] It is well known that, for normal dipole–dipole interactions in solution, the distance dependence of the attractive potential energy is 1/r6. In the case of hydrogen-bonding interactions, this is very complicated. The dependence varies from 1/r2 to 1/rx.[34] For XB, the only work, to the best of our knowledge, is from our recent publication. We found that the attractive XB energy between C6F5Cl and DMSO was (1/r3.1), whereas that between C6F5H and DMSO was (1/r3.3). Merz and Riley investigated the interaction energy of several halobenzene–formaldehyde complexes as a function of the halogen–oxygen separation distance.[32] It was found that the halogen-bond interaction energies increased as the size of the halogen atom increased. However, the distance-dependence equation was not provided. Herein, we examined the interaction energy between C6F5Br/C6F5I and DMSO and fitted the data to a function of Lennard–Jones form, E = a/rm + b/rn, and compared the results with those of the interaction pairs between C6F5Cl/C6F5H and DMSO. The dependence of the attractive term varied from 1/r3.1 to 1/r4.0, whereas that of the repulsive term was 1/r8.5 for all interaction pairs.
2. Results and Discussion 2.1. ATR and Excess IR Spectra Analysis First, we focus on the vs(C¢D) of [D6]DMSO, v(C¢Br) of C6F5Br, and v(C¢I) of C6F5I. The partial IR spectra of pure [D6]DMSO, DMSO, and the two benzene derivatives are presented in Figure 1. In Figure 1 A, the two bands centered at n˜ = 2124 and 2250 cm¢1 are attributed to the vs(C¢D) and vas(C¢D) of [D6]DMSO, respectively.[35] In Figure 1 B, the bands at n˜ 834 and 805 cm¢1 are attributed to v(C¢Br) in C6F5Br and v(C¢I) in C5F5I, respectively.[36] For [D6]DMSO, the band at n˜ 820 cm¢1 corresponds to the methyl rocking vibration.[35] To remove the influence of overlap, we used C6F5Br¢DMSO and C6F5I¢DMSO systems to analyze v(C¢Br) and v(C¢I). 2.1.1. C¢Br and C¢I Stretching Vibration in C6F5Br and C6F5I The IR and excess spectra in the v(C¢Br) and v(C¢I) regions of the C6F5Br¢DMSO and C6F5I¢DMSO systems are shown in Figures 2 and 3, respectively. A few features can be seen readily in the original IR spectra (Figures 2 A and 3 A) with increasing concentration of DMSO. First, v(C¢Br) and v(C¢I) gradually ChemPhysChem 2015, 16, 2594 – 2601
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Figure 1. Partial IR spectra of pure C6F5Br, C6F5I, DMSO, and [D6]DMSO in selected wavelength regions.
move to lower wavenumbers and the absorbance decreases. The wavenumber shifts of v(C¢Br) and v(C¢I) are shown in Figure S1 in the Supporting Information. The band positions of n(C¢Br) and v(C¢I) are redshifted by 1.1 and 3.5 cm¢1, respectively, which implies that the two single bonds are influenced by XBs and the halogen-bonding strength in the latter is greater. Second, the excess spectra in Figures 2 B and 3 B contain positive and negative bands and the positions of the positive bands are nearly fixed. The negative bands are at higher wavenumbers, whereas the positive bands are at lower wavenumbers compared with v(C¢Br) and v(C¢I) in pure C6F6Br and C6F5I, respectively. In the C6F5Br¢DMSO system, if the molar fraction of DMSO is less than 0.5, there are two (or even more) negative bands and one positive band. If the molar fraction of DMSO is greater than 0.5, the negative band is apparently singular (Figure 2 B). In the C6F5I¢DMSO system, if the molar fraction of DMSO is less than 0.3, there are also two negative bands and one positive band. If the molar fraction of DMSO is greater than 0.3, the negative bands look like singular ones (Figure 3 B). In an excess spectrum, the negative and positive bands represent the possible disappearing/increasing and appearing/increasing forms of the relevant species in the mixture compared with the pure liquid. In the lower concentration region of DMSO, there are two negative bands in v(C¢Br) and
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Figure 2. ATR-FTIR (A) and excess IR spectra (B) of the C6F5Br¢DMSO system in the range of v(C¢Br). From top to bottom: the molar fraction of DMSO increases from 0 to 1 in A) with an increment of about 0.1. The dashed and dash–dotted lines depict spectra of pure C6F5Br and DMSO in A). The inset model in A) is a sketch of C6F5Br self-association.
v(C¢I), which indicates that the amounts of the two species in the mixture decrease compared with those in the pure liquid. At higher concentrations of DMSO, the phenomenon is less clear. However, it is noted that the negative band is located between the two negative bands at lower concentrations of DMSO, which indicates that the broad, negative band at higher concentrations of DMSO is not a single band and may be composed of the two former negative bands. The two negative bands in the excess spectra indicate that there are at least two species in solutions of pure C6F6Br and C6F5I. The two species may be the C6F5Br/C6F5I monomer and C6F5Br/C6F5I selfassociated complex. Take C6F5Br self-association for example: the combination of C6F5Br monomers is through p–p stacking. A sketch of C6F5Br self-association is shown as the inset in Figure 2 A. The wavenumbers of v(C¢Br) and v(C¢I) in C6F5Br and C5F5I self-associates are lower than those of the C6F5Br and C6F5I monomers. When either of the solutions was diluted with DMSO, the amount of self-associates decreased to form C6F5Br¢DMSO and C6F5I¢DMSO complexes. The two negative bands in the excess spectra indicate the decrease in C6F5Br/ C6F5I monomer and C6F5Br/C6F5I self-associate, whereas the positive band over the entire concentration range indicates ChemPhysChem 2015, 16, 2594 – 2601
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Figure 3. ATR-FTIR (A) and excess IR spectra (B) of the C6F5I¢DMSO system in the range of v(C¢I). From top to bottom: the molar fraction of DMSO increases from 0 to 1 in A) with an increment of about 0.1. The dashed and dash–dotted lines depict spectra of pure C6F5I and DMSO in A).
the appearance of C6F5Br/C6F5I¢DMSO complexes. The wavenumber shifts of appearing species are lower than those of disappearing species. These results have been confirmed by quantum chemical calculations (see Section 2.3). An interesting feature in Figures 2 B and 3 B is the relative intensity of the negative bands when the molar fraction of DMSO is about 0.1. As indicated by the arrows in Figure 2 B, the intensity of the negative band in the higher wavenumber region is stronger than that in the lower wavenumber region. In Figure 3 B, however, the relative intensities of the two negative bands are opposite. This may be attributed to stronger interactions between C6F5I and DMSO than those between C6F5Br and DMSO. Thus, it is easier for DMSO to break apart the C6F5I self-associate, resulting in the stronger negative intensity of the band representing the self-associate (Figure 3 B). For the predominant breakup of the self-associate of C6F5Br, it requires a higher fraction of DMSO (0.2 and above). 2.1.2. vs(C¢D) in [D6]DMSO Shown in Figure 4 are the ATR-FTIR and respective excess spectra of vs(C¢D) in the C6F5Br¢[D6]DMSO and C6F5I¢[D6]DMSO systems over the entire concentration range. Two clear features
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Articles ple, at this concentration, the area of the negative band in the excess spectra of the C6F5Br¢[D6]DMSO system is barely seen, whereas that of C6F5I¢[D6]DMSO is quite pronounced. In combination with our former work,[18] for the C6F5H¢[D6]DMSO and ClC6F4H¢[D6]DMSO systems, the relative intensities of the negative bands are weaker than those of the C6F5I¢[D6]DMSO system. For the C6F5Cl¢[D6]DMSO system, the negative band does not appear until a molar fraction, x(C6F5Cl), of about 0.4. This indicates that C6F5I is a more powerful benzene derivative to break apart the DMSO self-associates, and the possible sequential order of the interaction strength of the five systems is as follows: C6F5I¢DMSO > C6F5H¢DMSO ClC6F4H¢DMSO > C6F5Br¢DMSO > C6F5Cl¢DMSO. This conclusion is supported by the quantum chemical calculation reported in Section 2.3.1. 2.2.
19
F NMR Spectroscopy Analysis
19
Figure 4. ATR-FTIR (A, B) and excess IR (C, D) spectra of the C6F5Br¢ [D6]DMSO (A, C) and C6F5I¢[D6]DMSO (B, D) systems in the range of vs(C¢D). From top to bottom: the molar fraction of [D6]DMSO decreases from 1 to 0 in A) and B) with a decrement of about 0.1. The dashed and dash–dotted lines depict spectra of pure [D6]DMSO, C6F5Br, and C6F5I.
can be seen readily in the original IR spectra (Figure 4 A and B): the absorbance of vs(C¢D) decreases monotonically and the bands gradually shift to higher wavenumber with increasing benzene derivatives. The wavenumber shifts of vs(C¢D) are shown in Figure S2 in the Supporting Information. Excess IR spectra provide another angle to observe molecular interactions between DMSO and the halogen-containing benzene derivatives, and the results are shown in Figure 4 C and D. General features of the two excess spectra are positive bands at higher wavenumbers and negative bands at lower wavenumbers in the two systems. With increasing molar fraction of benzene derivatives, the ratio of the absolute area of the negative band to the positive band increases gradually. The positions of positive and negative bands are nearly fixed. This means that DMSO molecules may have formed fixed interaction complexes in the solution. They are attributed to the absorptions of the [D6]DMSO molecules complexing with C6F5Br/C6F5I and those in self-associates. Another feature in Figure 4 C and D is that the areas of the negative bands in the excess spectra are different for the two systems. Taking the molar fraction of x(C6F5X) 0.1 as an examChemPhysChem 2015, 16, 2594 – 2601
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F NMR spectroscopy measurements of pure C6F5Br and C6F5I, and a series of C6F5Br¢DMSO, and C6F5I¢DMSO mixtures, were performed at 298 K, and the changes to the chemical shift of individual fluorine atoms during the dilution process were evaluated. Assignments of the 19F NMR spectroscopy signals to the fluorine atoms in C6F5Br and C6F5I are displayed in Figure 5. The concentration dependences of the spectra are displayed in Figure 6 A and B, and the dependences of chemical shift variations (Dd) of the two benzene derivatives on the molar fraction of DMSO are displayed in Figure S3 in Supporting Information. As observed, all chemical shift values become more and more negative with increasing concentration of DMSO, which means that the fluorine atoms gradually move upfield during the dilution process. This can be explained by the formation of halogen-bonding complexes because it is known that, upon forming XBs, electrons transfer from acceptors to donors, which causes upfield movement of the chemical shifts of the groups in the donors.[37, 38] Detailed charge gains of the fluorine atoms in the two benzene derivatives upon forming XB with DMSO have been calculated by the quantum chemical method, by using a 1:1 ratio in vacuum. The results are marked on the molecules in Figure 5. The negative values indi-
Figure 5. 19F NMR spectroscopy assignments of neat C6F5Br and C6F5I.
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Articles directional than XB. This is in agreement with conclusions reported in the literature.[40] Table 1 lists selected geometrical and vibrational parameters of the investigated complexes. For a better comparison, the corresponding data of the C6F5H¢DMSO and C6F5Cl¢DMSO systems is also listed.[18] As can be seen from these results, the C¢ Br and C¢I bonds in C6F5Br and C6F5I, respectively, are elongated upon complexation. The elongations are concomitant with a decrease in the stretching frequencies (redshift) and an increase in the IR intensities of the respective bonds. These are in agreement with the experimental results discussed in previous sections.
Table 1. Selected bond lengths (r in æ), vibrational frequencies (n in cm¢1), and IR intensities (I in km mol¢1) of the optimized complexes C6F5X¢DMSO (X = H, Cl, Br, or I). Data in parentheses represent changes in bond length (r), vibrational frequency (n), and intensity (I) upon complexation.
Complex
r(C¢X)
v(C¢X)
I(C¢X)
C6F5I¢DMSO
2.096 (0.016) 1.877 (0.004) 1.711 (0.002) 1.091 (0.007)
809.99 (¢1.81) 839.24 (¢1.71) 887.43 (¢0.14) 3157.67 (¢98.47)
72.83 (12.90) 109.79 (10.75) 151.14 (4.74) 318.13 (310.9)
C6F5Br¢DMSO C6F5Cl¢DMSO18
Figure 6. The molar fraction dependences of the 19F NMR spectroscopy signals in the C6F5Br¢DMSO (A) and C6F5I¢DMSO (B) systems.
C6F5H¢DMSO18
cate that all fluorine atoms in the two benzene derivatives gain charge upon forming XBs with DMSO.
2.3. Quantum Chemical Calculations 2.3.1. Structures The optimized geometries of individual molecules of DMSO, C6F5Br, and C6F5I, as well as the C6F5Br¢DMSO and C6F5I¢DMSO complexes, were studied. The sum of van der Waals atomic radii of bromine and oxygen (3.35 æ) and that of iodine and oxygen (3.55 æ) were used as critical values to determine the presence of a halogen bond.[39] The inset in Figure 7 shows the optimized geometries of the two complexes. The corresponding interaction energies of the complexes are displayed below the structures. The O¢X separations demonstrate that XBs form in both cases. Based on the interaction energies shown in Figure 7 and previous work,[18] we can easily obtain the sequential order of the interaction strength: C6F5I¢DMSO > C6F5H¢DMSO > C6F5Br¢DMSO > C6F5Cl¢DMSO. The strength of HB is between those of I¢XB and Br¢XB. Regarding the directionality of the halogen bonds, as shown in Figure 7, the optimized C¢Br···O and C¢I···O angles are all 1808. For C¢Cl···O angles,[18] they are 179 and 1768, which are very close to 1808. This reflects the characteristic directionality of XBs. In comparison, the optimized C¢H···O angles, as in the complexes of ClC6F4H¢DMSO and C6F5H¢DMSO, are 167 and 1638, which are much less than 1808 and show that HB are less ChemPhysChem 2015, 16, 2594 – 2601
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2.3.2. Angle Dependence of Energy To gain a deeper understanding of the directionality of halogen and hydrogen bonds, the dependences of the energies of XB and HB on the angle C¢X¢O (X = I, Br, Cl, and H) were studied. Single-point energy changes of the four complexes (C6F5H¢DMSO, C6F5Cl¢DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO)
Figure 7. Distance dependences of the interaction energies of the C6F5Cl¢ DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO complexes through halogen-bonding interactions. The inset shows the optimized geometries and corresponding interaction energies of the complexes. XB bonds are denoted by dashed lines, and the corresponding XB distances [æ] and bond angles [8] are labeled.
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Articles interactive nature changes from repulsive to attractive. There is a position with the lowest energy in each case that corresponds to the most stable optimized geometry shown in the inset of Figure 8. The data can be fitted to a function of Lennard–Jones form: E = a/rm¢b/rn. For Br¢XB, m = 8.5 and n = 3.2, and for I¢XB, m = 8.5 and n = 4.0. Combined with previous work,[18] the attractive energy of HB decays slightly slower than that of the I¢XB, although it is quicker than those of Br¢XB and Cl¢XB. All repulsive energies of HB and XB decay at 1/r8.5. 2.3.4. NBO Analysis Charge distributions of the most stable halogen-bonded complexes and the respective single molecules were examined by means of natural bond orbital (NBO) analysis. Some of the results are listed in Table 2. To compare HB and XB more clearly, the corresponding results for C6F5H¢DMSO and C6F5Cl¢DMSO are also listed in Table 2.[18]
Figure 8. XB and HB angle dependences of the interaction energies of the complexes C6F5H¢DMSO, C6F5Cl¢DMSO, C6F5Br¢DMSO, and C6F5I¢DMSO.
in their most stable forms as functions of the bond angle were calculated, and the results are shown in Figure 8 with the corresponding angle shown in the inset. The optimized positions are taken as the reference point (08 in Figure 8). The angle change is expressed by the angle differences between the most stable optimized structure and the analytical structures. Thus, with decreasing angle from the optimized position, the energies increase monotonically, which indicates that the interaction strength gradually reduces. More importantly, the angle dependences of XB or HB reflect the robustness of directionality. Thus, a decrease of 208 in the XB angle of C6F5I¢DMSO needs an energy compensation of about 6.7 kJ mol¢1. For the HB in C6F5H¢DMSO, however, the energy compensation is only about 1.6 kJ mol¢1. It is clear, therefore, that the directionality of the XB angles in C6F5X¢DMSO are very robust (ca. 1808). On the other hand, the directionality of the HB in C6F5H¢DMSO can be perturbed easily; a deviation of 138 from the straightline arrangement of the bond angle in C6F5H¢DMSO is observed.[18] Thus, the curves in Figure 8 demonstrate that the directionality of the three XBs follow the sequential order of I¢XB > Br¢ XB > Cl¢XB. In addition, all XBs have better directionality than that of HB. This can be explained by the fact that the positive electrostatic potential of a halogen atom is on the outermost portion of the atom, centered around the intersection of its surface with the C¢X axis (s hole), whereas the positive electrostatic potential of a hydrogen atom covers a much larger area of the atomic surface.[18] 2.3.3. Distance Dependence of Energy To study the dependence of XB energies on bond length, single-point energies of the two complexes in their most stable forms as functions of the Br···O and I···O separations were calculated, and the results are shown in Figure 7. The corresponding result for C6F5Cl¢DMSO is also drawn in the Figure 7.[18] As observed, with increasing separation, the energies change from positive to negative, which indicates that the ChemPhysChem 2015, 16, 2594 – 2601
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Table 2. Charges (q in e) on selected atoms and groups and charge change (CC) of the entire DMSO molecule. Data in parentheses represent CC values of the corresponding atoms or groups upon complexion. Negative values indicate increases in charge density. Complex
Methyl
X
O
CC
C6F5I¢DMSO
¢0.199 (0.038) ¢0.211 (0.026) ¢0.215 (0.022) ¢0.202 (0.035)
0.345 (0.055) 0.241 (0.043) 0.133 (0.034) 0.288 (0.046)
¢0.985 (¢0.031) ¢0.970 (¢0.016) ¢0.966 (¢0.012) ¢0.979 (¢0.025)
0.021
C6F5Br¢DMSO C6F5Cl¢DMSO18 C6F5H¢DMSO18
0.011 0.009 0.018
As observed from the results in Table 2, halogen-bonding donor molecules gain charges (negative charge) upon complexation, with markedly larger gain in the case of C6F5I¢ DMSO. Combined with former work,[18] the charge gains follow the sequential order of C6F5I¢DMSO > C6F5H¢DMSO > C6F5Br¢ DMSO > C6F5Cl¢DMSO. Clearly, stronger interactions result in the XB or HB donors having more charge. The halogen atoms, however, show a decrease in electron density, similar to a hydrogen atom when it acts as the HB donor. The oxygen atoms in the five complexes all gain charge upon complexation. In addition, the methyl groups of DMSO in all five complexes lose negative charge upon complexation. Stronger interactions result in greater charge loss from the methyl groups. Altogether, there is charge transfer from the hydrogen/halogen-bond acceptors (DMSO) to the donors. Although the methyl groups of DMSO do not participate directly in the formation of the halogen/hydrogen bonds, they donate electrons to the S=O group and contribute positively to the stability of the complexes. This is in agreement with our earlier work on the back groups.[18–20]
3. Conclusions By employing quantum chemical calculation and spectroscopic techniques, and taking several halogen-substituted benzene
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Articles derivatives, we performed a comparative study on XBs and HBs. Several conclusions could be drawn: 1) XBs were more directional than HBs. The optimized conformations of the interaction complexes in vacuum showed that Br¢XB and I¢XB were strict in this regard and Cl¢XB only slightly deviated from 1808. The angle dependences of the interaction energies further demonstrated that the directional order was I¢XB > Br¢ XB > Cl¢XB > HB. 2) Upon complexation, the C¢H and C¢X groups in benzene derivatives were all elongated, concomitant with decreases in the corresponding stretching frequencies (redshift) and increases in the IR intensities. 3) The interaction energy of the three XBs decreased with decreasing mass of the halogen atoms in the benzene derivatives. The interaction energy of HB was between those of Br¢XB and I¢XB. 4) The separation distance dependence of the interaction energies were described in the form of a Lennard–Jones equation. The attractive energies depended on 1/r3.3 for HB, and 1/r4.0, 1/r3.2, and 1/r3.1 for I¢XB, Br¢XB, and Cl¢XB, respectively; all repulsive energies of XB and HB decayed with 1/r8.5. Thus, stronger attractive interactions resulted in faster energy decays. This was true for HB and the three XBs. 5) Upon the formation of HB or XB, electrons transferred from DMSO to the benzene derivatives. Although the methyl groups of DMSO did not participate directly in the formation of HB/XB, they donated negative charge to the S=O group and contributed positively to the stability of the complexes and were not inert. Comparing the four interaction pairs, we found that stronger interactions made the methyl group lose more charge. These comparison studies on the interactions between fluorobenzene derivatives and DMSO provided in-depth information to understand the differences between XB and HB interactions in solution.
Experimental Section Materials C6F5Br (> 98 %) and C6F5I (> 98 %) were purchased from J&K. [D6]DMSO (> 99.8 % of deuterium) was purchased from Cambridge Isotopes Laboratories (USA). DMSO (> 99.5 %) was obtained from Beijing Chemical Plant (Beijing, P.R. China).
FTIR Spectroscopy FTIR spectra over the range from n˜ = 4000 to 650 cm¢1 were collected at room temperature ( 25 8C) by using a Nicolet 5700 FTIR spectrometer, equipped with a deuterated triglycine sulfate (DTGS) detector. The ATR cells, made of trapezoidal ZnSe and Ge crystals with incident angles of 45 and 608, corresponding to 12 and 7 reflections, respectively, were employed. The Ge crystal, with fewer numbers of reflections, and thus, shorter effective light path, was used to examine the strong stretching bands of C¢Br and C¢I. The ZnSe crystal was used to examine the weak stretching bands of C¢ D. All spectra were recorded with a resolution of 2 cm¢1, 16 parallel scans, and a zero filling factor of 2. For each sample, three parallel measurements were performed. The refractive indices of solutions were measured with an Abbe refractometer at 25 8C. The formula suggested by Hansen[41] was used to do the ATR corrections. ChemPhysChem 2015, 16, 2594 – 2601
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Excess IR Spectroscopy The theory of excess IR spectroscopy has been described in detail elsewhere.[19, 20] Briefly, an excess IR spectrum is defined as the difference between the spectrum of a real solution and that of the respective ideal solution under identical conditions. The working equation for calculating the excess IR spectrum is given by Equation (1): eE ¼
¨ ¦ A ¢ x1 e*1 þ x2 e*2 dðC1 þ C2 Þ
ð1Þ
in which A is the absorbance of the mixture; d is the light path length; Ci, xi, and e*i are the molarity, molar fraction, and molar absorption coefficient, respectively, of component i in its pure state. Matlab 7.0 (Math Works Inc., Natick, MA) was used to manipulate spectral data pretreatment, namely, the subtraction, truncation, and baseline correction, to calculate excess IR spectra and to integrate the bands in the excess IR spectra.
NMR Spectroscopy Measurements 19
F NMR spectroscopy is a sensitive indicator to provide intermolecular interaction information involving fluorine-substituted molecules, owing to its large chemical shift range and high intrinsic sensitivity (85 % of 1H NMR spectroscopy). The 19F NMR spectroscopy measurements were performed on a JEOL JNM-ECA 600 NMR (600 MHz) spectrometer at 298 K. Pure CF3COOH was used as an external standard to avoid its influence on the chemicals in the binary mixtures.
Quantum Chemical Calculations All computations were performed with the Gaussian 03 program.[42] We used the second-order many-body perturbation theory (MP2) with mixed basis sets to perform quantum chemical computations. The molecular energy, geometry, vibrational frequency, vibrational intensity, and natural population analysis (NPA) charge of all isolated single molecules and their complexes were fully optimized with the LANL2DZ[43] basis set for bromine and iodine atoms, and the aug-cc-pVDZ[44] basis set for hydrogen, carbon, fluorine, sulfur, and oxygen atoms. This allowed consideration of the relativistic effects in bromine and iodine atoms. The optimized geometries at the local energy minimum were ensured by the absence of imaginary vibrational frequency. The interaction energy was estimated as the difference between the total energy of the complex and the sum of total energies of the two individual molecules. Meanwhile, the basis set superposition error (BSSE) correction computed through the counterpoise method of Boys and Bernardi was estimated to obtain accurate interaction energies of the complexes.[45] The NPA charges were obtained by using the NBO approach.[46] To evaluate the distance dependence of the energy of XB, single-point energies were calculated by changing the distance between an electron donor (oxygen atom in DMSO) and an electron acceptor (bromine atom in C6F5Br and iodine atom in C6F5I).
Acknowledgements This work was supported by the Natural Science Foundation of China (grant nos. 21473099, 21273130).
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Articles Keywords: halogen bonds · hydrogen bonds · IR spectroscopy · structure–activity relationships · substituent effects [1] R. H. Crabtree, P. E. M. Siegbahn, O. Eisenstein, A. L. Rheingold, T. F. Koetzle, Acc. Chem. Res. 1996, 29, 348 – 354. [2] A. C. Legon, Phys. Chem. Chem. Phys. 2010, 12, 7736 – 7747. [3] A. Priimagi, G. Cavallo, P. Metrangolo, G. Resnati, Acc. Chem. Res. 2013, 46, 2686 – 2695. [4] M. Erd¦lyi, Chem. Soc. Rev. 2012, 41, 3547 – 3557. [5] A. C. Legon, Angew. Chem. Int. Ed. 1999, 38, 2686 – 2714; Angew. Chem. 1999, 111, 2850 – 2880. [6] T. Clark, M. Hennemann, J. S. Murray, P. Politzer, J. Mol. Model. 2007, 13, 291 – 296. [7] P. Politzer, P. Lane, M. C. Concha, Y. Ma, J. S. Murray, J. Mol. Model. 2007, 13, 305 – 311. [8] P. Politzer, J. S. Murraya, T. Clark, Phys. Chem. Chem. Phys. 2010, 12, 7748 – 7757. [9] P. Politzer, J. S. Murraya, ChemPhysChem 2013, 14, 278 – 294. [10] G. Trogdon, J. S. Murray, M. C. Concha, P. Politzer, J. Mol. Model. 2007, 13, 313 – 318. [11] P. Metrangolo, H. Neukirch, T. Pilati, G. Resnati, Acc. Chem. Res. 2005, 38, 386 – 395. [12] W. Z. Wang, N. B. Wong, W. X. Zheng, A. M. Tian, J. Phys. Chem. A 2004, 108, 1799 – 1805. [13] J. B. Cheng, R. Li, Q. Z. Li, B. Jing, Z. B. Liu, W. Z. Li, B. A. Gong, J. Z. Sun, J. Phys. Chem. A 2010, 114, 10320 – 10325. [14] J. P. M. Lommerse, A. J. Stone, R. Taylor, F. H. Allen, J. Am. Chem. Soc. 1996, 118, 3108 – 3116. [15] H. Umeyama, K. Morokuma, J. Am. Chem. Soc. 1977, 99, 1316 – 1332. [16] S. K. Noh, J. H. Jeon, W. J. Jang, H. Kim, S. H. Lee, M. W. Lee, J. Lee, S. Han, S. Kahng, ChemPhysChem 2013, 14, 1177 – 1181. [17] Q. Z. Li, Q. Q. Lin, W. Z. Li, J. B. Cheng, B. A. Gong, J. Z. Sun, ChemPhysChem 2008, 9, 2265 – 2269. [18] Y. Z. Zheng, N. N. Wang, Y. Zhou, Z. W. Yu, Phys. Chem. Chem. Phys. 2014, 16, 6946 – 6956. [19] Q. Z. Li, G. S. Wu, Z. W. Yu, J. Am. Chem. Soc. 2006, 128, 1438 – 1439. [20] Q. Z. Li, N. N. Wang, Q. Zhou, S. Q. Sun, Z. W. Yu, Appl. Spectrosc. 2008, 62, 166 – 170. [21] Y. Koga, F. Sebe, T. Minami, K. Otake, K. Saitow, K. Nishikawa, J. Phys. Chem. B 2009, 113, 11928 – 11935. [22] Y. Feng, D. Rainteau, C. Chachaty, Z. W. Yu, C. Wolf, P. J. Quinn, Biophys. J. 2004, 86, 2208 – 2217. [23] Y. Z. Zheng, N. N. Wang, J. J. Luo, Y. Zhou, Z. W. Yu, Phys. Chem. Chem. Phys. 2013, 15, 18055 – 18064. [24] N. N. Wang, Y. Wang, H. F. Cheng, Z. Tao, J. Wang, W. Z. Wu, RSC Adv. 2013, 3, 20237 – 20245.
ChemPhysChem 2015, 16, 2594 – 2601
www.chemphyschem.org
[25] J. Kiefer, M. M. Molina, K. Noack, ChemPhysChem 2012, 13, 1213 – 1220. [26] S. Corsetti, F. M. Zehentbauer, D. McGloin, J. Kiefer, Fuel 2015, 141, 136 – 142. [27] Z. P. Shields, J. S. Murray, P. Politzer, Int. J. Quantum Chem. 2010, 110, 2823 – 2832. [28] J. G. Hill, A. C. Legon, Phys. Chem. Chem. Phys. 2015, 17, 858 – 867. [29] R. Wilcken, M. O. Zimmermann, A. Lange, S. Zahn, B. Kirchner, F. M. Boeckler, J. Chem. Theory Comput. 2011, 7, 2307 – 2315. [30] D. Dolenc, B. Modec, New J. Chem. 2009, 33, 2344 – 2349. [31] H. S. El-Sheshtawy, B. S. Bassil, K. I. Assaf, U. Kortz, W. M. Nau, J. Am. Chem. Soc. 2012, 134, 19935 – 19941. [32] K. E. Riley, K. M. Merz, J. Phys. Chem. A 2007, 111, 1688 – 1694. [33] R. Li, Q. Z Li, J. B. Cheng, Z. B. Liu, W. Z. Li, ChemPhysChem 2011, 12, 2289 – 2295. [34] J. N. Israelachvili, Intermolecular and Surface Forces, 3rd ed., Academic Press, San Diego, 2011. [35] W. D. Horrocks, Jr., F. A. Cotton, Spectrochim. Acta 1961, 17, 134 – 147. [36] S. G. Frankiss, D. J. Harrison, Spectrochim. Acta 1975, 31, 1839 – 1864. [37] D. Hauchecorne, B. J. van der Veken, W. A. Herrebout, P. E. Hansen, Chem. Phys. 2011, 381, 5 – 10. [38] P. Metrangolo, W. Panzeri, F. Recupero, G. Resnati, J. Fluorine Chem. 2002, 114, 27 – 33. [39] L. Pauling, The Nature of the Chemical Bond, 3rd ed., Cornell University Press, New York, 1960. [40] C. Wang, D. Danovich, Y. Mo, S. Shaik, J. Chem. Theory Comput. 2014, 10, 3726 – 3737. [41] W. N. Hansen, Spectrochim. Acta 1965, 21, 815 – 833. [42] Gaussian 03, Revision B.03, M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A. Robb, J. R. Cheeseman, V. G. Zakrzewski, J. A. Montgomery, Jr., R. E. Stratmann, J. C. Burant, S. Dapprich, J. M. Millam, A. D. Daniels, K. N. Kudin, M. C. Strain, O. Farkas, J. Tomasi, V. Barone, M. Cossi, R. Cammi, B. Mennucci, C. Pomelli, C. Adamo, S. Clifford, J. Ochterski, A. Petersson, P. Y. G. Ayala, Q. Cui, K. Morokuma, D. K. Malick, A. D. Rabuck, K. Raghavachari, J. B. Foresman, J. Cioslowski, J. V. Ortiz, A. G. Baboul, B. B. Stefanov, G. Liu, A. Liashenko, P. Piskorz, I. Komaromi, R. Gomperts, R. L. Martin, D. J. Fox, T. Keith, M. A. Al-Laham, C. Y. Peng, A. Nanayakkara, C. Gonzalez, M. Challacombe, P. M. W. Gill, B. Johnson, W. Chen, M. W. Wong, J. L. Andres, C. Gonzalez, M. Head-Gordon, E. S. Replogle, J. A. Pople, Gaussian Inc., Pittsburgh, PA, 2003. [43] W. R. Wadt, P. J. Hay, J. Chem. Phys. 1985, 82, 270 – 283. [44] T. H. Dunning, J. Chem. Phys. 1989, 90, 1007 – 1023. [45] S. F. Boys, F. Bernardi, Mol. Phys. 1970, 19, 553 – 566. [46] A. E. Reed, L. A. Curtiss, F. Weinhold, Chem. Rev. 1988, 88, 899 – 926.
Received: April 18, 2015 Published online on June 26, 2015
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