Fluorine-Centered Halogen Bonding: A Factor in Recognition

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Fluorine-Centered Halogen Bonding: A Factor in Recognition Phenomena and Reactivity Published as part of the Crystal Growth & Design virtual special issue on Halogen Bonding in Crystal Engineering: Fundamentals and Applications Pierangelo Metrangolo,†,‡ Jane S. Murray,*,§ Tullio Pilati,† Peter Politzer,*,§ Giuseppe Resnati,†,‡ and Giancarlo Terraneo†,‡ †

NFMLab, Department of Chemistry, Materials and Chemical Engineering, “Giulio Natta”, Politecnico di Milano, Via L. Mancinelli 7, 20131 Milan, Italy ‡ CNST-IIT@POLIMI, Via Pascoli 70/3, 20133 Milan, Italy § CleveTheoComp, 1951 W. 26th Street, Suite 409, Cleveland, Ohio 44113, United States

bS Supporting Information ABSTRACT: Because of the anisotropies of their electronic charge distributions, many covalently bound halogen atoms have regions of positive electrostatic potential (positive σ-holes) on their outer portions. Through these, the halogens can interact attractively with negative sites. It has sometimes been questioned whether fluorine, the least polarizable halogen, can form halogen bonds, especially in solids. Here we present computational and crystallographic evidence demonstrating that it can indeed do so, in both the gaseous and the solid phases. We show computationally, through a series of examples, that fluorine can have positive σ-holes when linked to strongly electron-withdrawing residues and that it can interact with Lewis bases to form gas phase halogenbonded complexes. Through statistical analyses of data from the Cambridge Structural Database, we demonstrate that such fluorines do also halogen bond in the solid state, and we show several specific cases of this.

I. HALOGEN BONDING A covalently bonded halogen atom X in a molecule RX has an anisotropic electronic charge distribution;19 its radius is smaller along the extension of the RX bond than in the directions perpendicular to it. The region of lesser electronic density on the side of the atom opposite to the bond has been labeled a “σ-hole”.10 If the electronic density in the σ-hole is sufficiently low, then this region will have a positive electrostatic potential. For example, Figure 1 shows the electrostatic potential of orthobromochlorobenzene, computed on its molecular surface, which is taken to be the 0.001 au (electrons/bohr3) contour of its electronic density. The positive σ-holes of the bromine and chlorine are clearly visible. Their most positive values, designated as their VS,max, are 13.0 and 6.5 kcal/mol, respectively. In Table 1 are listed the computed VS,max for the halogen atoms X in a series of halides RX. Note that in some instances the VS,max are negative. This means that the anisotropy of the electronic charge surrounding the halogen is not sufficient to create a positive potential in the σ-hole. In such cases, there is still a σ-hole but it is negative, with VS,max < 0. The VS,max, whether positive or negative, are always on the extensions of the bonds RX and correspond to σ-holes. (The surface electrostatic potential around the lateral regions of a halogen is by definition less positive, or more negative, than the site of its VS,max.) The data in Table 1 show certain patterns, which have been observed in a great deal of earlier work.7,9,1115 For a given R, the r 2011 American Chemical Society

σ-hole potential of the halogen becomes more positive as its polarizability increases, that is, F < Cl < Br < I. In general, positive σ-holes are also promoted by the R portions of the molecules being more electron withdrawing; thus bromine has a higher VS,max in NCBr than in C6H5Br. A covalently bonded halogen can interact through a positive σ-hole with a negative site, such as the lone pair of a Lewis base.7,9,1119 This is called halogen bonding. Such interactions are highly directional, along the extensions of the covalent bonds to the halogens; when a halide RX forms a complex RX---B with a Lewis base B, the angle RXB is usually very close to 180° (barring secondary interactions). Halogen bonds are noncovalent interactions; the binding energies are in the same range as hydrogen bonds, and halogen bonding is indeed often competitive with hydrogen bonding.12,14,2023 It has been demonstrated that, for a given negative site, the strengths of halogen bonds correlate with the magnitudes of the σ-hole potentials.13,15,24 This emphasizes the largely electrostatic nature of these interactions. Table 2 gives the computed properties of some halogenbonded complexes RX---B. These data were taken from several sources and were obtained by different computational procedures. They are presented here simply as examples to illustrate general trends and possibilities. Received: July 13, 2011 Published: August 04, 2011 4238

dx.doi.org/10.1021/cg200888n | Cryst. Growth Des. 2011, 11, 4238–4246

Crystal Growth & Design

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Table 1. Comparison of Computed Electrostatic Potential Maxima (VS,max), in kcal/mol, on the 0.001 au Surfaces of the Halogen Atoms in Several Series of Halidesa molecule

VS,max

H3CF

21.0

H3CCl

1.5

H3CBr

5.9

C6H5F

13.3

C6H5Cl C6H5Br

Figure 1. Computed electrostatic potential on the 0.001 au molecular surface of ortho-bromochlorobenzene. The bromine and chlorine atoms are facing the viewer, at left and right, respectively. Color ranges, in kcal/ mol, are red, more positive than 10; yellow, from 10 to 5; green, from 5 to 0; blue, negative (less than 0). The most positive values, VS,max, are designated by black hemispheres; these are 13.0 (bromine) and 6.5 kcal/mol (chlorine). These correspond to the σ-holes on the bromine and the chlorine.

present work 14 14 present work

3.6 9.7

14 14 present work

F3CF

0.1

F3CCl

16.3

14

F3CBr

21.3

14

NCF

16.4

11

NCCl

34.9

11

NCBr

42.1

11

FF ClCl

13.8 23.8

14 11

BrBr

29.1

11

a

The values given are for the halogen indicated in bold. Calculations were at the B3PW91/6-31G(d,p) level.

Table 2. Computed Properties of Some Halogen-Bonded Complexes RX---Ba RXB angle

interaction energy, ΔE

H3COCl---NH3

180.0

4.3

12

H3COBr---NH3

180.0

7.0

12

NCF---NH3

176.6

1.9

present work

NCCl---NH3

179.8

6.5

present work

NCBr---NH3

179.8

8.0

present work

3.1 7.3

25 25

complex

The X---B separations are less than or approximately equal to the sums of the van der Waals radii of the respective atoms, which is consistent with these being noncovalent interactions. The angles RX---B are in the neighborhood of 180°, which is characteristic of halogen bonding. The largest deviations occur with H2CdO acting as the Lewis base; these are due to secondary interactions, as will be discussed later. The observed correlations between the strengths of halogen bonds and the magnitudes of the σ-hole potentials13,15,24 suggest that the interaction energies ΔE and the halogen VS,max reflect the same factors. This is confirmed by Table 2. If R and B are held constant, the interactions become stronger (more negative ΔE) in going from the lighter, less polarizable halogens to the heavier, more polarizable ones. The interactions are also strengthened by R being more electron withdrawing; compare, for instance, the ΔE for H3CI---NH3 and F3CI---NH3. A further influence upon ΔE is the nature of the Lewis base B. Thus, the complex BrCtCBr---NC4H4N (pyrazine) is more tightly bound than is BrCtCBr---NCCtCCN (Table 2) because the most negative electrostatic potentials on the nitrogen lone pairs in pyrazine are 31.3 kcal/mol vs. 23.4 kcal/mol for the nitrogen lone pairs in NCCtCCN. The final six entries in Table 2 demonstrate the possibility of forming chains of halogen-bonded molecules. An interesting feature is that when the atom with the σ-hole and the basic site are on the same molecule, successive interactions are increasingly stabilized, a “cooperativity effect”.9,11 For instance, the ΔE for NCBr---NCBr---NCBr is more than twice that for NCBr---NCBr. This can be attributed to polarization induced within each molecule by the interaction. When a bromine on molecule A interacts with the lone pair of a nitrogen on molecule B, electronic charge is polarized away from the bromine toward the

ref

H3CI---NH3 F3CI---NH3

>177 >177

ref

F3CCtCF---NH3

180.0

1.00

34

H3CCl---OdCH2

166.8

1.18

26

H3CBr---OdCH2

171.2

1.64

26

H3CI---OdCH2

172.9

2.32

26

BrCtCBr---NCCtCCN

180.0

2.8

11

BrCtCBr---NC4H4N

180.0

4.9

11

(pyrazine) NCF---NCF

180.0

1.3

11

NCF---NCF---NCF

180.0

2.9

11 11

NCBr---NCBr

180.0

4.9

NCBr---NCBr---NCBr

180.0

10.6

11

BrCtCBr---

180.0

5.5

11

NCCtCCN---BrCtCBr a

Distances are in angstroms, angles in degrees and energies in kcal/mol.

nitrogen on molecule A, making the latter more negative and able to interact more strongly with another bromine on molecule C. A positive halogen σ-hole is commonly surrounded by a region of negative electrostatic potential (Figure 1). This means that the same atom can interact electrostatically with both nucleophiles and electrophiles. This has indeed been observed via close contacts in crystals,2730 as well as computationally.31 4239

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Crystal Growth & Design Such dual reactivity cannot be explained in terms of single charges being assigned to atoms in molecules, by whatever method, a point that needs to be kept in mind in designing force fields.16,31 The importance of halogen bonding is increasingly being recognized in various areas, especially molecular biology, pharmacology, and crystal engineering.16,17,19,3236 It has been shown to be an important factor to consider and to exploit in designing new materials. When the electrostatic potential associated with a halogen σ-hole is negative, this generally precludes the possibility of halogen bonding, unless the electric field of the Lewis base is sufficiently strong to induce a positive σ-hole on the halogen.14 That this can happen is seen by noting that the σ-hole VS,max of the chlorine in H3CCl is near zero or weakly negative (Table 1), yet Riley and Hobza26 did obtain a ΔE of 1.18 kcal/mol for the interaction of H3CCl with OdCH2 (Table 2). Fluorine, the least polarizable halogen, often has a negative σ-hole and then does not halogen bond. In fact, it is sometimes stated or implied that fluorine never halogen bonds. This is incorrect. Table 1 includes examples of fluorines with both negative and positive σ-holes, and Table 2 cites computationally obtained σ-hole complexes involving fluorine. Legon has observed fluorine-centered halogen bonding experimentally, in the gas phase.37 Recent communications provide evidence for it occurring in the solid state as well.38,39 Our objective in this paper is to further substantiate that fluorine can and does halogen bond (albeit weakly) in both the gaseous and solid states. We wish thereby to direct attention to the reactivity and recognition possibilities afforded by fluorine-centered σ-hole interactions. Our approach shall draw upon computed electrostatic potentials and interaction energies, as well as surveys of crystallographic findings. Before proceeding in that direction, however, we would like to point out that covalently bonded atoms of Groups IVVI can also form positive σ-holes and interact through them with negative sites.14,19,40 The σ-holes are again on the extensions of the covalent bonds, so that Group IV, V, and VI atoms can have as many as four, three, and two positive σ-holes, respectively (or more, if hypervalent). The signs and magnitudes of the σ-hole potentials are governed by the same factors as for the halogens; accordingly, the first-row atoms (carbon, nitrogen, and oxygen) often do not have positive σ-holes.

II. COVALENTLY BONDED FLUORINE AS AN ELECTROPHILE Fluorine is generally regarded as the most electronegative atom, and in many organic molecules the electrostatic potential on its surface is entirely negative.10,13,15,39 It may therefore seem implausible that some covalently bound fluorines could halogen bond, which demonstrates electrophilic character. However the high electronegativity of the fluorine atom does not preclude some chemical groups from being even more electron-attracting. Experimentally based substituent constants, as in the recent compilation by Exner and B€ohm,41 show a number of groups, including NO2, CN, CF3, SO2CH3, etc., to be more electron-withdrawing than fluorine. Thus, in a molecule such as NCF, the fluorine could be expected to have some electrophilic tendencies, and Table 1 confirms that it does have a positive σ-hole. Among organic chemists, electrophilic fluorination has long been known.42 For instance, Barton et al. reported the addition of

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Figure 2. Computed electrostatic potential on the 0.001 au molecular surface of F3COF. Two views are shown: (a) The fluorine bonded to oxygen is at the right; one of the fluorines bonded to carbon is slightly to the left. (b) Two of the fluorines bonded to carbon are at the left. Color ranges, in kcal/mol, are red, more positive than 6; yellow, from 6 to 3; green, from 3 to 0; blue, negative (less than 0). The most positive values, VS,max, are designated by black hemispheres. The fluorine bonded to the oxygen has the most positive σ-hole, shown in panel (a).

fluoroxy compounds, for example, F3COF, to double bonds.42,43 The use of a variety of electrophilic fluorinating agents, including F2, (C6H5SO2)2NF, FClO3, and many others has been reviewed by Rozen44 and by Taylor et al.45 It seems reasonable to suggest that these fluorines, which are linked to strongly electron-withdrawing residues, may have positive σ-holes and that these could be responsible for the electrophilic behavior. Note the positive σ-hole on the OF fluorine of F3COF, Figure 2. The same explanation can be given for the reported formation of complexes between amines and perfluoroalkanes,46 and between F2 and NH3, HCN, and H2S.37 We and others have recently shown 4240



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Table 3. Computed Electrostatic Potential Maxima (VS,max) and Minima (VS,min), in kcal/mol, on the 0.001 au Surfaces of Covalently Bonded Fluorine Atoms in a Variety of Molecules. The Values Given Are for the Fluorine Indicated in Bold. Calculations Were at the B3PW91/6-31G(d,p) Level

a

Molecular surface of fluorine has totally positive electrostatic potential.

that the surface potential of F2 has positive σ-holes on both fluorines,14,34,39 which can interact with negative sites. Further examples of fluorine acting as an electrophile can be found in the crystal structures of compounds containing fluorine attached to a highly electron-attracting residue. A close intermolecular contact, RF---B, where B is a basic site on a neighboring molecule, can be interpreted as a σ-hole interaction if the F---B separation is less than (or very close to) the sum of the respective van der Waals radii and the RFB angle is in the vicinity of 180°. This combination has been observed on many occasions,30,38,4750 as shall be seen in Sections III and IV. (The RFB angle may deviate somewhat from 180° due to secondary interactions within the crystal lattice.)

III. COMPUTATIONAL ANALYSIS OF COVALENTLY BOUND FLUORINE AS AN ELECTROPHILE A. Procedure. In order to more fully characterize the possibility of covalently bound fluorine as an electrophile in σ-hole interactions, we have examined computationally the electrostatic potentials of a group of fluorine derivatives and looked at the interactions of some of them with Lewis bases. We will begin with a brief discussion of electrostatic potentials.

The nuclei and electrons of any molecule create an electrostatic potential V(r) at each point r in the surrounding space. It is given rigorously by the formula, Z ZA Fðr0 Þ dr0 ð1Þ V ðrÞ ¼ jr0 rj A jR A rj

∑

ZA is the charge on nucleus A, located at RA, and F(r) is the electronic density of the molecule. V(r) is a physical observable; it can be determined experimentally by diffraction techniques51,52 as well as computationally. Its sign in any region depends upon whether the positive contribution of the nuclei or the negative one of the electrons is dominant there. The electrostatic potential is a property of both fundamental and practical significance,53,54 the latter in particular as a guide to noncovalent interactions. For this purpose, it is commonly computed on a “surface” of the molecule, which is usually taken to be the 0.001 au (electrons/bohr3) contour of its electronic density, as suggested by Bader et al.55 Examples are shown in Figures 1 and 2. V(r) on the molecular surface is labeled VS(r); its locally most positive and most negative values, of which there may be several, are designated VS,max and VS,min. In this work, we have optimized molecular geometries and calculated V(r) with the B3PW91/6-31G(d,p) procedure. This has been used in previous studies,1113,18,31 and therefore allows comparisons of the fluorine VS,max to those of other halogen σ-holes. The optimized structures 4241



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and interaction energies ΔE of the complexes RF---B were obtained at the M06-2X/6-311G(3df,2p) level, the ΔE via eq 2, in which the energies are the minima at 0 K. ΔE ¼ EðR-F---BÞ EðR-FÞ EðBÞ

ð2Þ

The M06-2X density functional has been shown to be particularly well suited for treating weak interactions,56 such as those of present interest. The Gaussian 09 code57 was utilized for geometry optimizations, wave

Figure 3. Computed electrostatic potential on the 0.001 au molecular surface of FOF. One of the fluorines is facing the viewer, and the second is pointing to the right. Color ranges, in kcal/mol, are red, more positive than 6; yellow, from 6 to 3; green, from 3 to 0; blue, negative (less than 0). The most positive values, VS,max, are designated by black hemispheres.

functions and energies, and the Wave Function Analysis Surface Analysis Suite58 for V(r). The large basis set being used should minimize basis set superposition error,59 and this was accordingly not considered.

B. Survey of Covalently-Bound Fluorine-Centered σ-Hole Potentials. In Table 3 is a series of fluorine-containing molecules for

which we have computed the electrostatic potentials VS(r) on the 0.001 au molecular surfaces. Some of these have been used as electrophilic fluorinating agents.44,45 For each molecule, we list whatever VS,max and VS,min (most positive and most negative potentials) are found on the indicated fluorine. (In some instances, there is a positive or negative region on the fluorine surface that blends with that of a neighboring atom, and there is no VS,max or VS,min on the fluorine itself.) As in Table 1, the VS,max in Table 3, whether positive or negative, are associated with σ-holes and are on the extensions of the covalent bonds to the fluorines. Positive VS,max tend to occur when the R portion of RF is strongly electron withdrawing. The most positive fluorine VS,max in Table 3 are for NCCtCF and NCF; in these two molecules, as in several others, the entire surface of the fluorine is positive. However, there are a number of examples of fluorines with positive σ-holes that are surrounded by negative potentials, for instance, the fluorines bonded to the oxygens in F3COF (Figure 2a) and in F2O (Figure 3). These illustrate the point made earlier, that it can be quite misleading to assign a single positive or negative charge to an atom in a molecule.16,31 As will be seen, such fluorines can interact electrostatically through both their positive and their negative regions. The effect upon the σ-hole potential of a fluorine that is produced by the atom directly bonded to it can be seen, for instance, in F3C-OF and in FOCF2OF. The fluorines linked to the oxygens have significantly more positive σ-holes than do the fluorines on the carbons (Table 3 and Figure 2). Table 3 shows O2NF to be an exception to the usual tendency for electron-withdrawing R to strengthen σ-hole potentials. What is found instead in FNO2 is a positive “π-hole” above and below the nitrogen.60

Table 4. Computed M06-2X/6-311G(3df,2p) Interaction Energies, ΔE, F---B Separations and XFB Angles for Selected Gas Phase Interactions of Fluorine-Containing Molecules with Nitrogen Bases or Formaldehyde, or in a Self-Interacting Dimer. Energies and VS,max Are in kcal/mol, Distances in Angstroms and Angles in Degrees.

a

Reference 65. b This is the VS,max prior to interaction. 4242



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Figure 5. Ball-and-stick model of FOF---OdCH2 complex. The lower fluorine is interacting electrostatically through its positive σ-hole with a lone pair of the oxygen, while the upper hydrogen is interacting similarly with the negative sides of both fluorines.

Figure 4. Computed electrostatic potential on the 0.001 au molecular surface of formaldehyde, OdCH2. The oxygen is in the foreground, the hydrogens toward the back. Color ranges, in kcal/mol, are red, more positive than 15; yellow, from 15 to 0; green, from 0 to 20; blue, more negative than 20. The positions of the atoms are shown as gray spheres within the surface. The most negative values, VS,min, on the oxygen surface are designated by light blue hemispheres. They are at angles of 37° with respect to the 2-fold symmetry axis of the molecule. It is through this π-hole that occur the N---O noncovalent interactions between NO2 groups, both inter- and intramolecular, that have been observed crystallographically61,62 and computationally.60,6264 Table 3 contains, as a separate category, some molecules that are known to form networks in the solid state, through short intermolecular contacts; others in this group are candidates for doing so. An example is shown below, as structure 1.30 Note that it involves fluorinecentered σ-hole bonding, manifested in short F---O and F---F contacts and H---F hydrogen bonding. The F---F interactions are between the positive σ-hole on one fluorine and the negative lateral side of another, illustrating the point made above. Such networks, linked through a fluorine-centered σ-hole and other interactions, will be discussed in Section IV.

C. Some Fluorine-Centered σ-Hole Complexes. Table 4 presents computed properties of a series of complexes RF---B involving halogen bonding by fluorine to primarily NH3 or OdCH2. The ΔE values show most of the interactions to be relatively weak. The F---B separations are less than or in the vicinity of the sums of the respective van der Waals radii,65 which are included in the table. (It should be kept

in mind that van der Waals radii are simply approximate guides to noncovalent interactions and do not represent rigorous limits.) Most of the RF---B angles are close to 180°, but some of those in which OdCH2 is the base deviate quite considerably, for example, FOF---OdCH2. The oxygen of OdCH2 has two lone pairs, giving rise to two VS,min. These are located at angles of 37° with respect to the 2-fold symmetry axis (Figure 4). The fluorine-centered σ-hole interacts with one of the lone pairs, which allows one of the hydrogens of OdCH2 to participate in secondary interactions. For example, in FO-F---OdCH2, this hydrogen interacts attractively with the lateral sides of both fluorines, which are negative (Figure 3). Thus, this complex actually involves three interactions, as shown in Figure 5. The H---Flower and H---Fupper separations are 2.53 and 2.63 Å, respectively; the sum of the hydrogen and fluorine van der Waals radii is 2.67 Å.65 The two hydrogen bonds cause the OF---O interaction to deviate from linearity. Secondary interactions analogous to those in Figure 5 are likely to also be occurring to others of the complexes in Table 4, and they are found extensively in solids; see, for example, structure 1 and Chopra et al.30,38 It might accordingly be questioned whether fluorine-centered halogen bonding is actually taking place in these systems, or are the observed fluorine short contacts simply the result of other noncovalent interactions, especially in solids. This is not an issue in NCF---NCF and NCF---NC F---NCF (Table 2), in which linearity precludes secondary interactions, nor in NCF---NH3 (Table 4) or F3CCtCF---NH3 (Table 2), in which the hydrogens are well-removed from the nitrogen of NCF and the terminal fluorines of F3CCtCF. However, these complexes are known only computationally, which corresponds to the gas phase. Chopra and Guru Row have addressed the reality of fluorine-centered halogen bonds in some depth,38 and concluded that they do exist in solids as well as in the gas phase. The last entry in Table 4 shows an F---O interaction in a known molecular solid.66 The original report mentioned only O---C intermolecular interactions, but the structure in the Cambridge Structural Database (CSD) also reveals short F---O contacts of 2.946 Å. Our gas phase calculation shows that the F---O interaction alone gives a ΔE of 1.1 kcal/mol, with an F---O separation of 2.90 Å, similar to that found in the crystal.

IV. FLUORINE-CENTERED HALOGEN BONDING IN THE SOLID STATE Analyses of the CSD have proven useful in assessing trends in crystal structures and in providing indications of noncovalent interactions. For instance, as far as halogen bonding is concerned, the histogram of occurrences of CI---O short contacts (namely, interatomic distances shorter than the sum of the van der Waals radii of iodine and oxygen) versus CI---O contact angles shows a well-defined peak (median 165°, Supporting Information, Figure S1.1). This directionality is a consequence of the anisotropic 4243



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distribution of the electronic density around the covalently bound iodine, which results in a σ-hole on the outer side of the iodine that is marked enough to produce a tendency for interaction with an oxygen, a typical nucleophile, along the extension of the CI covalent bond. High-resolution X-ray diffraction studies67 have proven experimentally that the charge density distribution around a covalently bound fluorine atom is aspherical but much less so than around chlorine which, in turn, is less so than around bromine and iodine. The VS,max of the σ-holes, and the directionalities of the interactions with negative sites, tend to increase as the halogen anisotropy Table 5. Statistics from CSD on Short Contacts Involving Carbon-Bound Halogens entry

a

interaction

number of

median value of

structures

interaction angle (deg)

1

CF---OXa

937

130

2

CCl---OX

1778

159

3

CBr---OX

1024

163

4

CI---OX

525

165

5

CF---OX

28

139

6

CCl---OX

56

162

7

CBr---OX

29

167

8 9

CI---OX CF---Met+

22 82

170 123b

10

CCl---Met+

34

103b

11

SO2CF2F---OX

112

138

X is defined as any atom. b Statistics for the interaction of a metal cation with bromine and iodine are not given due to the small number of structures available in CSD.

becomes greater. The median value of the CX---O angle, when a neutral oxygen atom interacts with a neutral halogen, decreases monotonically from iodine to fluorine (Table 5, entries 14, Supporting Information, Figures S1.1S1.4). This trend confirms that fluorine is the halogen with the weakest positive σ-holes, if any. The absolute value of the median CF---O angle (130°) suggests that most fluoroorganics either do not have a positive σ-hole or that it is too weak to favor, in the solid, the interaction with a nucleophile on the extension of the covalent CF bond. A careful analysis of small differences in specific populations of the CSD and a focus on trends in particular subsets give clear indications that the anisotropic electronic distribution around a covalently bound fluorine affects the geometry of fluorinecentered interactions and identifies cases where such a distribution may function as structure-controlling. Effective mapping of the asymmetric electronic distributions around halogens can be performed by examining the interactions of charged species with the halogens. Interactions with oxyanions (Table 5, entries 58, Supporting Information, Figures S1.5S1.8) are much more linear than with neutral oxygen atoms, while metal cations (Table 5, entries 9, 10, Supporting Information, Figures S1.9S1.10) are attracted to the negative lateral sides of the halogen and interact nearly orthogonally to the CX bonds; the respective angles for CF compounds are 139° and 123°. These trends and absolute values of directionalities confirm some degrees of asymmetry in the electronic distributions around carbon-bound fluorine atoms. This asymmetry consists of a negative belt and a positive, or less negative, cap; these affect interactions in crystalline solids. The interaction with a neutral oxygen atom shows a linear preference, quite weak, only with the subset of fluorine atoms linked to strong electron-withdrawing residues, namely, only moieties wherein the overall electronic density on the fluorine

Figure 6. Ball-and-stick representations of 0D system formed by 1-ethyl-3-methyl-4-nitroimidazolium trifluoromethanesulfonate (A), CSD code GEDLEH, 1D infinite chain (view along a) formed by bis(1-(difluoro-oxo-λ6-sulfanylidene)-2-fluoro-2-oxoethyl)-mercury (B), CSD code VIHNEF and of the supramolecular ribbon (view along b) formed by 1,3,3-tetrakis(trifluoromethanesulfonyl)propane (C), CSD code SIZPIB. The digits close to the XB (dashed lines) are the angle formed by covalent and noncovalent bonds around the fluorine atom. Color code: gray, carbon; yellow, sulfur; red, oxygen; blue, nitrogen; light green, fluorine; purple, mercury. In (A) hydrogen atoms are omitted for clarity. 4244



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manifest itself when no stronger competing interactions determine the crystal packing, since these may preferentially dominate and prevent weak fluorine-centered halogen bonds from forming.

’ ASSOCIATED CONTENT

bS

Supporting Information. Analyses of the Cambridge Structure Database (CSD) and selected structures from the CSD that meet the standard geometric requirements for fluorine functioning as an XB donor. This material is available free of charge via the Internet at http://pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected] (P.P.), [email protected] (J.S.M.).

’ ACKNOWLEDGMENT P.M., G.R., and G.T. thank Fondazione Cariplo (project 2150, New-Generation Fluorinated Materials as Smart Reporter Agents in 19F MRI and project 2010-1351, Development of a Technology Platform between the South and the North of Europe: Exchange Research Program between Politecnico di Milano and VTTTechnical Research center of Finland (S2N)). Figure 7. Ball-and-stick representations of the 2D grid formed by 1,3, 5-trifluoro-1,1,3,5,5-pentanitropentane (A), CSD code CUVFAA and (μ2-1-fluoro-2-(trifluoromethyl)but-1-en-1-yl-4-ylidyne)-nonacarbonyl-triiron (B), CSD code FAFYUH. The digits close to the XB (dashed lines) are the angle formed by the covalent and noncovalent bonds around the fluorine atom. Color code as in Figure 6, light brown, iron. Hydrogen atoms are omitted for clarity. A CF(NO2)2 moiety and some CO moieties have been removed for clarity in (A) and (B), respectively.

has been depleted (Table 5, entry 11, Supporting Information, Figure S1.11). A positive σ-hole, and the resulting halogen bond, may thus influence the overall crystal packing in systems in which fluorine is bonded to strongly electronegative groups. An analysis of structures in the CSD that meet the standard geometric requirements for fluorine functioning as a halogen bond donor reveals that this is most likely to take place when no stronger competing interactions govern the crystal packing; the latter may prevent fluorine-centered halogen bonds, which are weak interactions, from occurring. Examples of zero-,68 one-,69 and twodimensional70 networks are presented in Figures 6 and 7, and many other structures show short and linear contacts (Supporting Information, Figures S2.1S2.4), demonstrating the ability of fluorine to act as a halogen bond donor.30,71 Disorder is frequently present in crystals containing fluorinated residues, but those cited above show no disorder, so that generalizations drawn from contact distances and angles are reliable.

V. SUMMARY While the presence of a positive σ-hole on the outer side of a covalently bound fluorine in fluorocarbons is relatively uncommon, we have presented computational and crystallographic evidence that it can and does occur, in the solid state as well as the gaseous phase, in compounds in which the fluorine atoms are linked to strong electron-withdrawing groups. The ability of fluorine to act as a halogen bond donor in solids is most likely to

’ REFERENCES (1) Stevens, E. D. Mol. Phys. 1979, 37, 27–45. (2) Nyburg, S. C.; Wong-Ng, W. Proc. R. Soc. London A 1979, 367, 29–45. (3) Ikuta, S. J. Mol. Struct. (Theochem) 1990, 205, 191–201. (4) Pedireddi, V. R.; Reddy, D. S.; Goud, B. S.; Craig, D. C.; Rae, A. D.; Desiraju, G. R. J. Chem. Soc., Perkin Trans. 2 1994, 2353–2360. (5) Price, S. L.; Stone, A. J.; Lucas, J.; Rowland, R. S.; Thornley, A. E. J. Am. Chem. Soc. 1994, 116, 4910–4918. (6) Tsirelson, V. G.; Zou, P. F.; Tang, T.-H.; Bader, R. F. W. Acta Crystallogr. 1995, A51, 143–153. (7) Lommerse, J. P. M.; Stone, A. J.; Taylor, R.; Allen, F. H. J. Am. Chem. Soc. 1996, 118, 3108–3116. (8) Torii, H. J. Chem. Phys. 2003, 119, 2192–2198. (9) Grabowski, S. J.; Bilewicz, E. Chem. Phys. Lett. 2006, 427, 51–55. (10) Clark, T.; Hennemann, M.; Murray, J. S.; Politzer, P. J. Mol. Model. 2007, 13, 291–296. (11) Politzer, P.; Murray, J. S.; Concha, M. C. J. Mol. Model. 2007, 13, 643–650. (12) Politzer, P.; Murray, J. S.; Lane, P. Int. J. Quantum Chem. 2007, 107, 3046–3052. (13) Riley, K. E.; Murray, J. S.; Politzer, P.; Concha, M. C.; Hobza, P. J. Chem. Theor. Comput. 2009, 5, 155–163. (14) Politzer, P.; Murray, J. S.; Clark, T. Phys. Chem. Chem. Phys. 2010, 12, 7748–7757. ezac, J.; Sola, R. J.; (15) Riley, K .E.; Murray, J. S.; Fanfrlík, J.; R Concha, M. C.; Ramos, F. M.; Politzer, P. J. Mol. Model. 2011, DOI: 10.1007/s00894-011-1015-6. (16) Auffinger, P.; Hays, F. A.; Westhof, W.; Shing Ho, P. Proc. Nat. Acad. Sci. U. S. A. 2004, 101, 16789–16794. (17) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Acc. Chem. Res. 2005, 38, 386–395. (18) Politzer, P.; Lane, P.; Concha, M. C.; Ma, Y.; Murray, J. S. J. Mol. Model. 2007, 13, 305–311. (19) Murray, J. S.; Riley, K. E.; Politzer, P.; Clark, T. Aust. J. Chem. 2010, 63, 1598–1607. (20) Di Paolo, T.; Sandorfy, C. Chem. Phys. Lett. 1974, 26, 466–469. (21) Di Paolo, T.; Sandorfy, C. Can. J. Chem. 1974, 52, 3612–3622. 4245


Crystal Growth & Design (22) Corradi, E.; Meille, S. V.; Messina, M. T.; Metrangolo, P.; Resnati, G. Angew. Chem., Int. Ed. 2000, 39, 1782–1786. (23) Alkorta, I.; Blanco, F.; Solimannejad, M.; Elguero, J. J. Phys. Chem. A 2008, 112, 10856–10863. (24) Shields, Z.; Murray, J. S.; Politzer, P. Int. J. Quantum Chem. 2010, 110, 2823–2832. (25) Valerio, G.; Raos, G.; Meille, S. W.; Metrangolo, P.; Resnati, G. J. Phys. Chem. A 2000, 104, 1617–1620. (26) Riley, K. E.; Hobza, P. J. Chem. Theor. Comput. 2008, 4, 232–242. (27) Murray-Rust, P.; Motherwell, W. D. S. J. Am. Chem. Soc. 1979, 101, 4374–4376. (28) Murray-Rust, P.; Stallings, W. C.; Monti, C. T.; Preston, R. K.; Glusker, J. P. J. Am. Chem. Soc. 1983, 105, 3206–3214. (29) Ramasubbu, N.; Parthasarathy, P.; Murray-Rust, P. J. Am. Chem. Soc. 1986, 108, 4308–4314. (30) Chopra, D.; Thiruvenkatam, V.; Manjunath, S. G.; Guru Row, T. N. Cryst. Growth Des. 2007, 7, 868–874. (31) Politzer, P.; Murray, J. S.; Concha, M. C. J. Mol. Model. 2008, 14, 659–665. (32) Voth, A. R.; Hays, F. A.; Shing Ho, P. Proc. Nat. Acad. Sci. 2007, 104, 6188–6193. (33) Metrangolo, P.; Resnati, G.; Pilati, T.; Liantonio, R.; Meyer, F. J. Polym. Sci. A 2007, 45, 1–15. (34) Lu, Y.-X.; Zou, J.-W.; Yu, Q.-S.; Jiang, Y.-J.; Zhao, W.-N. Chem. Phys. Lett. 2007, 449, 6–10. (35) Metrangolo, P.; Resnati, G.; Pilati, T.; Biella, S. In Halogen Bonding: Fundamentals and Applications; Metrangolo, P.; Resnati, G., Eds.; Springer: Berlin, 2008; pp 105136. (36) Rissanen, K. CrystEngComm 2008, 10, 1107–1113. (37) Legon, A. C. Angew. Chem., Int. Ed. 1999, 38, 2686–2714. (38) Chopra, D.; Guru Row, T. N. CrystEngComm 2011, 13, 2175–2186. (39) Metrangolo, P.; Murray, J. S.; Pilati, T.; Politzer, P.; Resnati, G.; Terraneo, G. CrystEngComm 2011, DOI: 10.1039/c1ce05554b. (40) Politzer, P.; Murray, J. S. In Practical Aspects of Computational Chemistry; Leszczynski, J.; Shukla, M. K., Eds.; Springer: Amsterdam, 2009; pp 149167. (41) Exner, O.; B€ohm, S. J. Phys. Org. Chem. 2006, 19, 393–401. (42) Barton, D. H. R.; Godinho, L. S.; Hesse, R. H.; Pechet, M. M. Chem. Commun. 1968, 804–806and references cited. (43) Barton, D. H. R; Danks, L. J.; Ganguly, A. K.; Hesse, R. H.; Tarzia, G.; Pechet, M. M. Chem. Commun. 1969, 227–228. (44) Rozen, S. Chem. Rev. 1996, 96, 1717–1736. (45) Taylor, S. D.; Kotoris, C. C.; Hum, G. I. Tetradedron 1999, 55, 12431–12477and references cited. (46) Burdeniuc, J.; Sanford, M.; Crabtree, R. H. J. Fluorine Chem. 1998, 91, 49–54. (47) Hamor, T. A.; Watkin, D. J. J. Chem. Soc., Perkin Trans. 2 1974, 140–146. (48) Polyakov, O. G.; Ivanova, S. M.; Gaudinski, C. M.; Miller, S. M.; Anderson, O. P.; Strauss, S. H. Organometallics 1999, 18, 3769–3771. (49) Deck, P. A.; McCauley, B. D.; Slebodnick, C. J. Organomet. Chem. 2006, 691, 1973–1983. (50) Takahashi, A.; Yanai, H.; Taguchi, T. Chem. Commun. 2008, 2385–2385. (51) Stewart, R. F. Chem. Phys. Lett. 1979, 65, 335–342. (52) Politzer, P.; Truhlar, D. G., Eds. Chemical Applications of Atomic and Molecular Electrostatic Potentials; Plenum Press: New York, 1981. (53) Politzer, P.; Murray, J. S. Theor. Chem. Acc. 2002, 108, 134–142. (54) Murray, J. S.; Politzer, P. Wiley Interdisc. Rev., Comp. Mol. Sci. 2011, 1, 153–163. (55) Bader, R. F. W.; Carroll, M. T.; Cheeseman, J. R.; Chang, C. J. Am. Chem. Soc. 1987, 109, 7968–7979. (56) Zhao, Y.; Truhlar, D. G. Theor. Chem. Acc. 2008, 120, 215–241. (57) Gaussian 09, Revision A.1; Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.;

ARTICLE

Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, Jr., J. A.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, N. J.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, € Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. A. D.; Farkas, O.; Gaussian, Inc.: Wallingford, CT, 2009. (58) Bulat, F. A.; Toro-Labbe, A.; Brinck, T.; Murray, J. S.; Politzer, P. J. Mol. Model. 2010, 16, 1679–1691. (59) Grimme, S. J. Comput. Chem. 2006, 27, 1787–1799. (60) Murray, J. S.; Lane, P.; Clark, T.; Riley, K. E.; Politzer, P. J. Mol. Model. 2011, DOI: 10.1007/s00894-011-1089-1. (61) Wozniak, K.; He, H.; Klinowski, J.; Jones, W.; Grech, E. J. Phys. Chem. 1994, 98, 13755–13765. (62) G€obel, M.; Tchitchanov, B. H.; Murray, J. S.; Politzer, P.; Klap€otke, T. M. Nat. Chem. 2009, 1, 229–235. (63) Murray, J. S.; Lane, P.; G€obel, M.; Klap€otke, T. M.; Politzer, P. J. Chem. Phys. 2009, 130, 104304(16). (64) Macaveiu, L.; G€obel, M.; Klap€otke, T. M.; Murray, J. S.; Politzer, P. Struct. Chem. 2010, 21, 139–146. (65) Bondi, A. J. Phys. Chem. 1964, 68, 441–451. (66) Della Vedova, C. O.; Downs, A. J.; Novikov, V. P.; Oberhammer, H.; Parsons, S.; Romano, R. M.; Zawadski, A. Inorg. Chem. 2004, 43, 4064–4071. (67) Hathwar, V. R.; Guru Row, T. N. Cryst. Growth Des. 2011, 11, 1338–1346. (68) Katritzky, A. R.; Yang, H.; Zhang, D.; Kirichenko, K.; Smiglak, M.; Holbrey, J. D.; Reichert, W. M.; Rogers, R. D. New J. Chem. 2006, 30, 349–358. (69) Bittner, J.; Seppelt, K. Chem. Ber. 1991, 124, 87. Takahashi, A.; Yanai, H.; Taguchi, T. Chem. Commun. 2008, 2385–2387. (70) Atovmyan, L. O.; Golovina, N. I.; Eremenko, L. T.; Zhitomirskaya, N. G.; Oreshko, G. V.; Fadeev, M. A. Izv. Akad. Nauk SSSR, Ser. Khim. (Russ.) (Russ. Chem. Bull.) 1984, 33, 543; Lentz, D.; Willemsen, S. J. Organomet. Chem. 2002, 641, 215–219. (71) Hamor, T. A.; Watkin, D. J. J. Chem. Soc., Perkin Trans. 2 1974, 140–146. (CSD code BODFNB)Polyakov, O. G.; Ivanova, S. M.; Gaudinski, C. M.; Miller, S. M.; Anderson, O. P.; Strauss, S. H. Organometallics 1999, 18, 3769–3771. (CSD code QAPHOE)Nockemann, P.; Thijs, B.; Parac-Vogt, T. N.; Van Hecke, K.; Van Meervelt, L.; Tinant, B.; Hartenbach, I.; Schleid, T.; Ngan, V. T.; Nguyen., M. T.; Binnemans, K. Inorg. Chem. 2008, 47, 9987–9999. (CSD code TOJREQ)Nockemann, P.; Thijs, B.; Pittois, S.; Thoen, J.; Glorieux, C.; Van Hecke, K.; Van Meervelt, L.; Kirchner, B.; Binnemans, K. J. Phys. Chem. B 2006, 110, 20978–20992. (CSD code JESSUL)Platonov, V. E.; Haas, A.; Schelvis, M.; Lieb, M.; Dvornikova, K. V.; Osina, O. I.; Gatilov, Yu. V. J. Fluorine Chem. 2001, 109, 131–139(CSD code XIHLUV).

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