Competition between hydrogen bond and halogen bond

PAPER

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Competition between hydrogen bond and halogen bond in complexes of formaldehyde with hypohalous acids Qingzhong Li,*a Xisen Xu,a Tao Liu,b Bo Jing,a Wenzuo Li,a Jianbo Cheng,a Baoan Gonga and Jiazhong Suna Received 14th December 2009, Accepted 17th March 2010 First published as an Advance Article on the web 11th May 2010 DOI: 10.1039/b926355a An ab initio study of the complexes formed by hypohalous acids (HOX, X = F, Cl and Br) with formaldehyde has been carried out at the MP2/aug-cc-pVTZ computational level. Two minima complexes are found, one with an H O contact and the other one with an X O contact. The former is more stable than the latter, and the strength difference between them decreases as the size of the X atom increases. The associated HO and XO bonds undergo a bond lengthening and red shift, whereas a blue shift was observed in the bond of the hypohalous acid not involved in the interaction. The interaction strength and properties in both complexes are analyzed with atoms in molecules (AIM) and natural bond orbital (NBO) theories. The energy decomposition analyses indicate that the contribution from the electrostatic interaction energy is larger in the hydrogen-bonded complexes than that in the halogen-bonded complexes.

1. Introduction Intermolecular interactions play an important role in supramolecular assembly, crystal packing, reaction selectivity, and drug–receptor interactions.1–3 Hydrogen bonding is, without doubt, the most important one among the various types of intermolecular interactions and has attracted much attention since the concept of hydrogen bond was put forward.4 Recently, a growing number of experimental and theoretical evidences confirm that the interactions, named as halogen bonds, play a crucial role in biochemistry and medicinal chemistry.5–7 For example, halogen bonds are often involved in protein–ligand interactions that are either biologically detrimental or beneficial. Halogen bonds8 are formed between halogen-containing molecules and Lewis bases. Initially, halogen-bonded complexes were taken as charge-transfer complexes because the charge transfer interaction is considered to play a dominant role in the complex formation.9 By means of energy decomposition scheme, scientists have an insight into the nature of halogen bonds and think that the electrostatic interaction, polarization, charge transfer, and dispersion are all responsible for the stability of halogen-bonded complexes.10 Recently, scientists have appreciated a deeper recognition that halogen bonding involves the interaction of positive sigma-holes (positive regions of electrostatic potential on the extensions of the covalent bond to the halogen) on the halogen with negative sites on other molecules.8,11 It has been shown that halogen bonds are analogical to hydrogen bonds in many aspects.12,13 In a

The Laboratory of Theoretical and Computational Chemistry, Science and Engineering College of Chemistry and Biology, Yantai University, Yantai 264005, P. R. China. E-mail: [email protected]; Fax: (+086) 535 6902063; Tel: (+086) 535 6902063 b Department of Chemistry, Jining University, Qufu 273155, P. R. China

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strength, halogen bonds are comparable to hydrogen bonds and thus a competition may occur between them. A competition of sigma-hole bonding with hydrogen bonding has been studied with theoretical methods.14 As powerful oxidizing agents, hypohalous acids (HOX, X = F, Cl, Br and I) can participate in oxidation, epoxidation and hydroxylation processes.15 Hypohalous acids are unstable and readily form OX radicals in the atmosphere, thus they appear as reservoirs and sinks of OX radicals in atmospheric chemistry,16–18 imposing a great influence on the stratospheric ozone concentration. It is therefore not surprising that they have been extensively investigated both experimentally and theoretically.19–22 Berski et al.22 studied the bonding in hypohalous acids with the topological analysis of the electron localization function at the B3LYP and Hartree–Fock levels and concluded that the O–F bond is a covalent, polarized bond, whereas the bonding between O and Cl, Br and I atoms is of electron donor–acceptor-type with the halogen donating electron density to the valence shell of oxygen. Due to the high electronegativity and small size of the fluorine atom, the behavior of hypofluorous acid (HOF) is dramatically different from the heavier hypohalous acids which exhibit similar properties. A number of studies have been performed for the hydrogen-bonded clusters formed by the hypohalous acids with themselves23,24 or with other molecules such as ozone,25 hydroperoxy radical,26 nitrosyl hydride,27 and sulfur trioxide.28 As known, the HOX species (X = F, Cl, or Br) are easily formed in the atmosphere and the formaldehyde emission in the atmosphere is increasing in modern society. Therefore, the title complexes in the present study are of great interest in the field of atmospheric chemistry relevant to ecological aspects. The present work presents a detailed theoretical study of the stabilities, electronic structures, and vibrational frequencies of the complexes of H2CO HOX (X = F, Cl or Br). The properties and interaction nature have also been understood with natural bond orbital theory, atoms in molecules theory, Phys. Chem. Chem. Phys., 2010, 12, 6837–6843 | 6837

and energy decomposition analysis. To the best of our knowledge, neither theoretical nor experimental data regarding the information about the interaction between hypohalous acids and formaldehyde are available in the literature. In the absence of experimental information, a theoretical analysis of the existence of such complexes and their properties appears to be necessary for understanding the interaction mechanism between them.

2. Theoretical methods All calculations were performed using the Gaussian 03 system of codes.29 The geometries of the complexes and the respective monomers were optimized at the MP2(FC)/aug-cc-pVTZ level. This level of theory adequately describes hydrogen bonds and halogen bonds,30–32 and thus was used in this study. The harmonic vibrational frequency calculations were then performed to confirm the predicted structures are minima and evaluate the zero-point vibrational energy (ZVE) and thermal energy (TE). The interaction energy is calculated as the difference between the energy of the complex and the total energy of the monomers. The counterpoise (CP) procedure of Boys and Bernardi33 was used to correct the interaction energy for basis set superposition error (BSSE). The natural bond orbital (NBO) analyses were carried out using the NBO package34 included in the Gaussian 03 suite of programs. The atoms in molecules (AIM) calculations were performed using the AIM2000 program.35 The scheme developed by Ziegler and Rauk36,37 was applied to perform the energy decomposition analysis. In this approach, the total interaction energy between two fragments is decomposed into three terms: DEint = Eelst + EPauli + Eoi. The term Eelst is the ‘‘rigid’’ electrostatic interaction energy between the fragments, which is calculated from the wave functions of the separated fragments. The term EPauli is the Pauli repulsive energy (electrons with the same spin cannot occupy the same region in space) between the fragments. The orbital interaction energy (Eoi) is the interaction energy of the occupied orbitals on one fragment and unoccupied orbitals on another. The energy decomposition was performed at the PBE/QZ4P (small core) level with ADF program.38

3. Results and discussion 3.1

Geometrics

The molecular geometries of hypohalous acids in gas phase have been determined by means of microwave spectroscopy.39–41 At the MP2/aug-cc-pVTZ computational level, Alkorta et al.30 have investigated the geometrics of hypohalous acids and found that the obtained geometrics are in good agreement with the experimental values. In the present paper, we reoptimized the structures of three hypohalous acids (HOX, X = F, Cl and Br) at the same level and the obtained results support the above conclusion. For example, the calculated O–Cl bond length is 1.697 A˚ and the experimental one is 1.689 A˚.40 Furthermore, this level of theory adequately described the structures and properties of hydrogen- and halogen-boned complexes involved with hypohalous acids.30,31 6838 | Phys. Chem. Chem. Phys., 2010, 12, 6837–6843

Fig. 1 Molecular electrostatic potential of the HOX (X = F, Cl and Br) molecule at 0.027 a.u. isosurfaces calculated at the MP2/aug-cc-pVTZ level. The red color represents the minimal molecular electrostatic potential and the blue color denotes the maximal molecular electrostatic potential.

The electrostatic potential has been found to be an effective tool for analyzing and predicting noncovalent interactions.42–44 Fig. 1 displays the molecular electrostatic potentials of the three hypohalous acid molecules calculated at the MP2/ aug-cc-pVTZ level. Among the electronic characteristics of the isolated monomers it is worth mentioning the presence of two maxima in the molecular electrostatic potentials of HOCl and HOBr but one maximum in the molecular electrostatic potentials of HOF. In the former two molecules, the halogen surface is negative, but there is a positive region on the outermost portion of the halogen surface, where it intersects the O–X axis. For each complex between the hypohalous acid and formaldehyde, we thus only considered two main configurations, one through a hydrogen bonding and the other through an O X contact (Fig. 2), although other configurations may also be formed. All atoms in the six complexes are in a plane. The O atom of H2CO acts as the electron donor in both interactions. In the halogen-bonded complexes, the free H atom of HOX is distant from the H atom of H2CO, whereas the free X atom of HOX is close to the H atom of H2CO in the hydrogen-bonded complexes. The distance between the X atom of HOX and the H atom of H2CO in the hydrogen-bonded complexes shortens as the size of the X atom decreases. The minimum molecular electrostatic potential of the X atom in the hypohalous acids is 0.0289, 0.0163 and 0.0134 au for the F, Cl and Br atoms, respectively.30 This shows that there is a secondary interaction of CH X in the hydrogen-bonded complexes. Although this secondary interaction may be very weak, it is important in maintaining the configuration of the hydrogen-bonded complexes. The geometrical parameters in the six complexes are presented in Table 1. The hydrogen bond angle is a little larger than the halogen bond angle, and both of them increase as the size of the X atom increases. Taylor and Kennard45 surveyed 111 crystal structures including hydrogen bonds and proposed a van der Waals cutoff that the H Y distance is smaller than the sum of the corresponding atom van der Waals radii in most H Y hydrogen bonds. This sum is about 2.5, 2.8, 3.2 and 3.4 A˚ for the O atom with H, F, Cl and Br atoms, respectively. It is seen from Table 1 that the van der Waals cutoff is valid in most complexes except in the H2CO–FOH complex. The larger binding distance in this complex shows that there is no halogen bonding between H2CO and FOH, supporting the conclusion that F atom not favored to participate in formation of a halogen bond.11 The H2CO–FOH complex may be associated This journal is

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Table 2 Selected stretch frequencies (v/cm1) and their shifts (Dv) in the complexes calculated at the MP2/aug-cc-pVTZ level Complex

v(H–O)

Dv(H–O)

v(X–O)

Dv(X–O)

H2CO–FOH H2CO–HOF H2CO–ClOH H2CO–HOCl H2CO–BrOH H2CO–HOBr

3762 3518 3777 3522 3781 3525

1 245 1 253 10 246

971 984 744 774 640 674

9 4 21 9 23 11

mass of X atom. Interestingly, the ‘distant’ H–O and X–O bonds of the hypohalous acid not involved in the interaction undergo a blue shift in all complexes regardless of the change of the bond length. The distant blue shift of the X–O stretch is larger than that of the H–O stretch, which is reverse to the change of the above red shift in the complexes. This distant blue shift has also been observed in complexes of hypohalous acids with nitrosyl hydride27 and carbon monoxide.31 Fig. 2 Optimized structures of the complexes between H2CO and HOX (X = F, Cl and Br) at the MP2/aug-cc-pVTZ level.

3.3 Interaction energies Table 1 Binding distance (R/A˚), bond lengths (r/A˚) and their changes (Dr/A˚), halogen and hydrogen bond angle (y/degree) in the complexes calculated at the MP2/aug-cc-pVTZ level Complex

R

r(X–O) Dr(X–O) r(H–O) Dr(H–O) y


2.958 1.802 2.702 1.803 2.652 1.812

1.430 1.433 1.706 1.692 1.836 1.816

0.002 0.004 0.009 0.005 0.014 0.007

0.970 0.983 0.968 0.981 0.968 0.981

0.000 0.013 0.000 0.013 0.001 0.012

162.4 167.6 170.5 172.3 172.1 172.7

simply through a van der Waals force. For the three hydrogenbonded complexes, the higher the electronegativity of the distant X atom in HOX, the shorter the binding distance. However, the reverse is found for the binding distance in the H2CO–XOH complexes although the radii of X atom increases. Upon complexation, the H–O and X–O bonds are lengthened. The F–O bond elongation in H2CO–FOH complex is the smallest, whereas the Br–O bond lengthening in the H2CO–BrOH complex is the largest. The elongation of H–O and X–O bonds has a relation with the binding distance. The shorter binding distance is accompanied with a larger elongation of H–O and X–O bonds. The free H–O and X–O bonds do not follow a simple change in formation of the complex. The free H–O bond in H2CO–FOH and H2CO–ClOH complexes almost shows no change, the free F–O bond in the H2CO–HOF complex displays a small elongation, while the other distant H–O and X–O bonds are shortened. 3.2

Frequency shifts

Table 2 presents the H–O and X–O stretch frequencies and their shifts in the complexes at the MP2/aug-cc-pVTZ level. The associated H–O and X–O bonds exhibit a red shift, which is consistent with the elongation of the bonds. Here it is mentioned that blue shifts are also observed in some halogen and hydrogen bonds.46 The red shift of the X–O stretch is much smaller than the H–O stretch red shift due to the heavier This journal is

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The interaction energies in the complexes are calculated at the MP2/aug-cc-pVTZ level. The results are given in Table 3. These energies are corrected with BSSE, ZVE and TE. The results are also presented in Table 3. The BSSE value corresponds to about 9–22% of the uncorrected interaction energy, the ZVE value amounts to about 19–42% of the uncorrected interaction energy, while the TE value equals about 28–140% of the uncorrected interaction energy. Clearly, the TE proportion is the largest, followed by the ZVE counterpart, and the BSSE contribution is the smallest. The effect of three terms on the O X contact is larger than that on the hydrogen bond. In the three O X contacts, the effect of all three terms increases as the X atom varies from F to Br. The three terms have the largest effect on the interaction energy of the H2CO–FOH complex. The interaction energy corrected with TE leads to a positive value for H2CO–FOH complex, indicating this molecular pair is unfavorable. In the following discussion, we only focus on the interaction energy corrected with BSSE for comparison with the results in literature. The interaction energy is 31.01, 30.47 and 29.26 kJ mol1 for H2CO–HOF, H2CO–HOCl and H2CO–HOBr complexes, respectively. The strength of hydrogen bond in the three complexes is consistent with the acidities of the hypohalous acids (1554 (HOF), 1395 (HOCl), 1354 (HOBr) kJ mol1 at 298 K).47 Of course, the nature of the halogen atom has relatively little impact on the strength of hydrogen bond. Hypofluorous acid is a slightly stronger acid than HF,47 thus the interaction energy of the hydrogen bond in the H2CO–HOF complex is larger than that in the H2CO–HF complex (4.85 kcal mol1 at the MP4(SDTQ)/ 6-311++G(2df,2pd)// MP2/6-311++G(2df,2pd) level).45 The interaction energy is 23.41 and 20.73 kJ mol1 for H2CO–HCl and H2CO–HBr complexes, respectively. The strength of the hydrogen bond in the complex of H2CO with the other hypohalous acids is also stronger than for the hydrohalic acids HX (X = Cl and Br) although the Phys. Chem. Chem. Phys., 2010, 12, 6837–6843 | 6839

Table 3 Uncorrected interaction energy (DE/kJ mol1), interaction energies corrected with BSSE (DECP/kJ mol1), ZVE (DEZVE/kJ mol1) and TE (DETE/kJ mol1), and deformation energy (DE/kcal mol1) in the complexes calculated at the MP2/aug-cc-pVTZ level Complexes

DE

BSSE

DECP

ZVE

DEZVE

TE

DETE

DE

DECCSDa


5.42 34.25 15.69 33.85 24.53 34.31

1.15 3.24 1.92 3.38 5.32 5.05

4.27 31.01 13.77 30.47 19.21 29.26

2.25 8.11 4.09 7.57 4.66 7.55

3.17 26.14 11.6 26.28 19.87 26.76

7.58 9.76 7.88 9.60 8.11 9.63

2.16 24.49 7.81 24.25 16.42 24.68

0.02 0.69 0.14 0.74 0.30 0.81

3.26 30.31 9.74 27.91 13.64 26.31

a

DECCSD was obtained with a single point calculation at the CCSD/aug-cc-pVTZ level on the MP2/aug-cc-pVTZ geometry.

hypohalous acids are weaker than the hydrohalic acids HX (X = Cl and Br).47 In order to estimate the effect of halogen atoms in the hydrogen-bonded complexes, the similar structure of the complex between H2CO and water has also been optimized at the same level. The obtained interaction energy (21.32 kJ mol1) clearly indicates the strong withdrawing effect of the halogen atom of the hypohalous acids that favors a stronger hydrogenbonding interaction. A similar result has also been reported by Alkorta.30 The nature of the halogen atom has a prominent influence on the strength of the O X contact. The interaction energy in the H2CO–FOH complex is the smallest (4.27 kJ mol1). The strength of halogen bond has a relation with the positive electrostatic potential on the extension of the X–O bond and the positive potential correlates with the polarizability of the halogen atom.11,48,49 The greater polarizability of the Br atom relative to F and Cl results in the larger strength of halogen bond. As the size of the X atom increases, the interaction energy of the halogen bonding between the hypohalous acid and H2CO is close to that of the hydrogen-bonding energy between them. When a mixed basis set (Lanl2DZ* has been used for the I atom and aug-cc-pVTZ for the remainder of the atoms) is adopted for the complexes of H2CO with IOH, the interaction energy of the halogen bond (16.05 kJ mol1) in the H2CO–IOH complex is close to that of the hydrogen bond (17.00 kJ mol1) in the H2CO–HOI complex. A similar result has also been observed in the complexes of hypohalous acids with three nitrogenated bases (NH3, N2 and NCH).30 For HOI and N2, the strength of halogen bond is even a little stronger than the hydrogen bond.30 This shows that the strength of halogen bonds is comparable with that of hydrogen bonds, thus there is a competition between hydrogen bonds and halogen bonds in some cases.50 We also consider the effect of deformation energy on the interaction energy. The deformation energy is the monomer deformation energy in the complex and is calculated to be the difference between the energy sum of the monomers in the complex and that in the isolated molecules. The result is also presented in Table 3. Evidently, the deformation energy is very small in both types of interactions. This indicates that the deformation of monomers in the complexes can be neglected in calculating the interaction energy. The deformation energy is larger in the hydrogen-bonded complexes and smallest in the H2CO–FOH complex. The magnitude of the deformation energy reflects the strength of the interaction. 6840 | Phys. Chem. Chem. Phys., 2010, 12, 6837–6843

Finally, we also calculated the interaction energies at the CCSD/aug-cc-pVTZ level, which were performed with a single-point energy calculation on the MP2/aug-cc-pVTZ geometries. The result is also given in Table 3. The interaction energies at the CCSD/aug-cc-pVTZ level are less negative than those at the MP2/aug-cc-pVTZ level, but the change sequence is the same for both methods. With the increase of X atom mass, the difference between the two levels is increased. 3.4 NBO analyses In formation of O HO hydrogen bond, an orbital interaction happens between the lone pair of the electron donor and the antibonding H–O orbital as anticipated. A similar interaction has been found in the halogen boned complexes. The stabilization energies due to these orbital interactions are given in Table 4. A large n(O) - s*(H–O) orbital interaction is observed in the H2CO–HOX complexes, which can be used to explain the elongation and red shift of the H–O bond.51 The n(O) - s*(X–O) orbital interaction in the H2CO–XOH complexes is relatively weak and thus the elongation and red shift of the X–O bond are also small. The nature of the halogen atom has a small effect on the n(O) - s*(H–O) orbital interaction, whereas it imposes a great impact on the n(O) - s*(X–O) orbital interaction. The n(O) - s*(Br–O) orbital interaction energy is two times larger than the n(O) - s*(Cl–O) counterpart, whereas the n(O) - s*(F–O) orbital interaction is negligible in the H2CO–FOH complex. On the basis of the change consistence of the orbital interaction energy with the interaction energy, we believe that the orbital interaction plays an important role in the complex conformation. In formation of hydrogen and halogen bonds, a charge transfer occurs from the electron donor to the electron acceptor. A similar charger transfer is observed in all the complexes except the H2CO–FOH complex (Table 4). In this complex, a small charge transfers from FOH to H2CO. It is reasonable if we consider the close contact between the F atom in FOH and the H atom in H2CO. The charge transfer in the hydrogenbonded complexes is larger than that in the halogen-bonded complexes. As the size of the X atom increases, the charge transfer decreases in the former, whereas it increases in the latter. For the Br complexes, the charge transfer is almost equal in both types of interactions. The change of charge transfer is consistent with the interaction energy, indicating that the charge transfer is of importance in formation of both types of interactions. This journal is

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Table 4 Stabilization energy (E/kcal mol1), charge transfer (CT/e), differences between NBO electron density (ED) in the complexes and the isolated HOX in XO and HO sigma bonding (Ds) and sigma antibonding (Ds*) orbitals in the complexes calculated at the HF/aug-cc-pVTZ wave functiona Complexes

E

CT

Ds(XO)

Ds*(XO)

Ds(HO)

Ds*(HO)


0.12 16.60 4.18 15.77 8.82 13.83

0.0019 0.0204 0.0061 0.0178 0.0154 0.0156

0.0008 0.0004 0.0029 0.0008 0.0031 0.0010

0.0002 0.0004 0.0105 0.0002 0.0201 0.0005

0.0001 0.0003 0.0001 0.0004 0.0001 0.0004

0.0001 0.0222 0.0001 0.0199 0.0001 0.0189

Stabilization energy is due to the n(O) - s*(X–O) and n(O) - s*(H–O) orbital interactions for the O X contact and O HO hydrogen bond, respectively. Charge transfer is the difference of H2CO charge in the complex with that in the isolated molecule. a

Investigating the electron density (ED) changes in the XO subsystems, we found an insignificant ED decrease in the associated bonding s(XO) orbital and a much larger increase of ED in the antibonding s*(XO) orbital in H2CO–XOH (X = Cl and Br) complexes. Both effects, but mainly the latter one, are responsible for the elongation and weakening of the X–O bond. A similar situation is found for the associated s(HO) and s*(HO) orbitals in H2CO–HOX (X = Cl and Br) complexes. For H2CO–FOH complex, the increase of ED in the s*(FO) orbital is very small, which is consistent with the very small orbital interaction (0.12 kcal mol1). For the H2CO–HOF complex, the ED changes in the associated s(HO) and s*(HO) orbitals both increase, but the latter has a much greater increase. The ED in the free bonding and antibonding orbitals suffers a very small change. 3.5

AIM analyses

The topological analysis of the electron density indicates the presence of intermolecular bond critical points (BCPs) in all the complexes (Fig. 3). One BCP is found in the halogenbonded complexes, whereas three BCPs are seen in the other complexes. The electron density and Laplacian at the main BCP are presented in Table 5. In all cases, these bond critical points have a small value of the electron density and positive Laplacian, which are an indication of the closed-shell characteristics of the interaction.52 The values of both terms are larger in the H2CO HOX complexes than in the H2CO XOH complexes, which is consistent with the interaction strength in both complexes. This relationship can be further confirmed in Fig. 4, where a linear relationship is found between the interaction energy and the electron density at the intermolecular BCP. Diagonalization of the Hessian of the electron density yields three eigenvalues (l1, l2 and l3). The ellipticity (e = (l1/l2) 1) measures the extent to which charge is preferentially accumulated, thus it provides a measure for the stability of bond.53 Substantial bond ellipticity means that the bond can easily be ruptured. The ellipticity of O F BCP in H2CO–FOH is very large, which supports this complex is very weak. Fig. 5 shows the relationship of the interaction energy with the ellipticity at the intermolecular BCP. Evidently, it shows that there are three different complexes: hydrogenbonded complexes (the ellipticity is smallest), halogen-bonded complexes (the ellipticity is smaller), and van der Waals complex (the ellipticity is largest). This journal is

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Fig. 3 Molecular graphs of the complexes between H2CO and HOX (X = F, Cl and Br) at the MP2/aug-cc-pVTZ computational level. Small red balls indicate the bond critical points.

Fig. 4 Interaction energy vs. electron density at the intermolecular BCP in all of the complexes calculated at the MP2/aug-cc-pVTZ computational level.

We also considered the change of electron density at the XO and HO BCPs in H2CO XOH (X = F, Cl and Br) complexes Phys. Chem. Chem. Phys., 2010, 12, 6837–6843 | 6841

Table 5 Electron density (r/au), corresponding Laplacian (r2r/au) and ellipticity (e) at the intermolecular bond critical point (BCP), changes of electron density (Dr/au) at the XO and HO BCPs in H2CO XOH (X = F, Cl and Br) complexes calculated at the MP2/aug-cc-pVTZ level 2

Complex

rinter

r rinter

einter

Dr(XO)

Dr(HO)


0.0053 0.0366 0.0159 0.0331 0.0229 0.0323

0.0288 0.1010 0.0698 0.1019 0.0975 0.1013

0.4355 0.0216 0.0561 0.0119 0.0810 0.0070

0.0009 0.0070 0.0039 0.0042 0.0057 0.0050

0.0005 0.0205 0.0009 0.0202 0.0012 0.0193

Table 6 Decomposition of interaction energies (in kcal mol1) in H2CO XOH (X = F, Cl and Br) complexes at the PBE/QZ4P(small core) level in ADF program Complex

DEinta

EPauli Eelst


3.20 32.09 14.51 30.47 16.76 24.57

4.26 48.81 31.80 48.39 53.81 49.44

a

4.73 49.81 25.39 48.28 41.92 47.68

Eoi

Eelst (%) Eoi (%)

2.73 31.09 20.92 30.59 28.65 26.33

63.40 61.57 54.83 61.21 59.40 64.42

36.60 38.43 45.17 38.79 40.60 35.58

DEint = Eelst + EPauli + Eoi.

4. Conclusions

Fig. 5 Interaction energy vs. ellipticity at the intermolecular BCP in all of the complexes calculated at the MP2/aug-cc-pVTZ computational level.

(Table 5). In all complexes, the electron density of the associated XO and HO BCPs decreases, resulting in a weakening of the bonds. In hydrogen- and halogen-bonded complexes, the electron density of the distant XO and HO BCPs increases, accompanied with an enhancement of the bonds, whereas it decreases in H2CO–FOH and H2CO–HOF complexes. 3.6

Energy decomposition analyses

In order to unveil the nature of the interaction in these complexes, we performed a fragment-based energy decomposition analysis for the interaction energies. They were decomposed into three parts: electrostatic interaction energy (Eelst), Pauli repulsion energy (EPauli), and orbital interaction energy (Eoi). The results are presented in Table 6. The Eelst and Eoi terms are negative, showing a positive contribution to the stability of the complexes, whereas the EPauli value is positive. For hydrogenbonded complexes, the magnitude of Eelst is close to EPauli, whereas for halogen-bonded complexes, the absolute value of Eelst is larger than EPauli. In all complexes, the contribution of Eelst to the interaction energy is larger than that from Eoi. The Eelst contribution in hydrogen-bonded complexes is larger than that in halogen-bonded complexes. In the conceptual Kohn–Sham framework, the Eoi term accounts basically for charge-transfer and donor–acceptor orbital interactions between the two fragments, and can be considered as a measure of the covalent character of the intermolecular bond54 The bigger contribution of the Eoi in halogen-bonded complexes indicates that the halogen bonding has a more covalent character. 6842 | Phys. Chem. Chem. Phys., 2010, 12, 6837–6843

Quantum chemical calculations have been performed to study the complexes formed between the hypohalous acids (HOX, X = F, Cl and Br) and formaldehyde. The H2CO–ClOH and H2CO–BrOH complexes are combined with a halogen bond, whereas the H2CO–FOH complex is associated with a van der Waals force. The energetic results indicate that the hydrogenbonded complexes are favored and the strength of halogen bond is close to that of hydrogen bond as the size of the X atom increases. The nature of the halogen atom has relatively little impact on the strength of the hydrogen bond, whereas it has a prominent influence on the strength of the O X contact. In all complexes, the associated HO and XO bonds are elongated and show a red shift, while the free bond in the hypohalous acids displays a blue shift regardless of the change of the bond length. Analysis of the electron density of the complexes indicates a linear relationship between electron density at the bond critical point and the interaction energy. Although the NBO analysis shows that the charge transfer and orbital interaction between the bonding orbital of H2CO and the antibonding orbital of the HOX molecule is important for the stabilization of the complexes, the energy decomposition analysis indicates that the electrostatic interaction plays a more important role in all complexes. We think that this work is beneficial for understand the mechanism of interaction between the hypohalous acids and formaldehyde, which can be confirmed with spectroscopic methods in the future.

Acknowledgements This work was supported by the National Natural Science Foundation of China (Grant No. 20973149). It was also supported in part by open project of state key laboratory of supramolecular structure and materials (SKLSSM200909) from Jilin University, China.

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