Fluorine–Fluorine Interactions in the Solid State - ACS Publications

ARTICLE pubs.acs.org/JPCA

Fluorine
Fluorine Interactions in the Solid State: An Experimental and Theoretical Study Robert J. Baker,†,* Paula E. Colavita,† Deirdre M. Murphy,† James A. Platts,‡ and John D. Wallis§ †

School of Chemistry, University of Dublin, Trinity College, Dublin 2, Ireland School of Chemistry, Main Building, Cardiff University, Park Place, Cardiff, CF10 3AT, U.K. § School of Science and Technology, Nottingham Trent University, Nottingham, NG11 8NS, U.K. ‡

bS Supporting Information ABSTRACT: The solid state structures of three compounds that contain a perfluorinated chain, CF3(CF2)5CH2CH(CH3)CO2H, CF3(CF2)5(CH2)4(CF2)5CF3 and {CF3(CF2)5CH2CH2}3PdO have been compared and a number of C
F 3 3 3 F
C and C
F 3 3 3 H
C interactions that are closer than the sum of the van der Waals radii have been identified. These interactions have been probed by a comprehensive computational chemistry investigation and the stabilizing energy between dimeric fragments was found to be 0.26
29.64 kcal/mol, depending on the type of interaction. An Atoms-in-Molecules (AIM) study has confirmed that specific C
F 3 3 3 F
C interactions are indeed present, and are not due simply to crystal packing. The weakly stabilizing nature of these interactions has been utilized in the physisorption of a selected number of compounds containing long chain perfluorinated ponytails onto a perfluorinated self-assembled monolayer, which has been characterized by IRRAS (Infrared Reflection Absorption Spectroscopy).

’ INTRODUCTION Noncovalent interactions such as hydrogen bonding,1 π
π stacking,2 anion
π stacking,3 cation
π stacking,4 lone pair
π interactions,5 or C
H 3 3 3 π stacking6 have become cornerstones of supramolecular chemistry.7 In contrast, noncovalent interactions between the halogens are not used significantly in the building of supramolecular architectures.8 The heavier halogens can undergo X 3 3 3 X interactions; however, F 3 3 3 F interactions are somewhat contentious as, according to Pauling’s principle,9 fluorine has a low polarizability, which means the attractive interatomic dispersion forces would be rather low.10 Indeed, in 1,2,3,4-tetrafluorobenzene there is a preference for C
H 3 3 3 F
C interactions over C
F 3 3 3 F
C interactions.11 One recent application that involves the use of C
F 3 3 3 H
C interactions as an integral part of a supramolecular array has been shown to improve charge carrier properties for organic thin-film transistors.12 Another notable use for noncovalent interactions has been in the use of Teflon to aid precipitation of highly fluorinated phosphines from solution in catalytic studies.13 Further synthetic studies have paved the way for substitutions of one or more C
H groups for a C
F group that profoundly affect enzyme
ligand binding affinities.14 Thus the field of organic fluorine chemistry is of significant interest. There are a large number of structures where the fluorous and nonfluorous domains of the molecule pack together, which are often ascribed to packing forces. Notwithstanding this, there are a r 2011 American Chemical Society

number of reports in the literature of C
F 3 3 3 F
C interactions (specifically, those that contribute a stabilization energy) that have been characterized both experimentally and theoretically. A structural analysis of C—X 3 3 3 X—C interactions reported in 1986 that there are two types of interaction as shown in Chart 1.15 It was stated that type I interactions are not stabilizing and formed by close packing only. Uracil derivatives whereby one C
H moiety has been replaced by a C
F group show substantially different solid state structures, with the C
F 3 3 3 F
C interaction being analyzed theoretically and shown to impart a stabilization of 1
4 kcal/mol. Interestingly, the same authors analyzed the CSD database and found that the majority of C
F 3 3 3 F
C interactions are of type I, suggesting that these are not necessarily due to close packing.16 These interactions have recently been reviewed,10,17 but some pertinent examples are included here (Chart 2). One polymorph of A shows type II C
F 3 3 3 F
C interactions in addition to C
H 3 3 3 O and C
H 3 3 3 π interactions.18 Charge density studies on this and a related compound, B, have shown these can be classed as well-defined, but weak interactions.19 Derivatives of 1,8-difluoronaphthalene have been studied computationally and shown that type II C
F 3 3 3 F
C interactions occur with an Received: October 18, 2011 Revised: November 23, 2011 Published: November 27, 2011 1435

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The Journal of Physical Chemistry A Chart 1. Classification of Halogen
Halogen Interactions

Chart 2

internuclear separation of 2.3
2.8 Å, and imparts up to 14 kcal/mol of local stability.20 Directing abilities have been seen in the complex [Ag 3 (ntb)2(CF 3 SO3 )3 ] (ntb = tris(2-benzimidazolylmethyl)amine) where a supramolecular structure of the anionic (CF3SO3)6 cluster is held together by twelve C
F 3 3 3 F
C interactions as well as N
H 3 3 3 O hydrogen bonds.21 Hexafluoroacetylacetone copper compounds have also shown C
F 3 3 3 F
C interactions that direct self-assembly.22 Therefore, it is clear that further investigations of closed shell C
F 3 3 3 F
C interactions are required to gain a fundamental understanding and enable them to be further utilized in supramolecular chemistry. We have an interest in the use of fluorinated compounds in applications to environmental chemistry and separation science. Recently we reported on the use of fluorinated phenols for the extraction of Cs+ and Sr2+ from water into toluene,23 and the use of highly fluorinated phosphine oxides,24 Rf3PdO and Rf2P(dO)C2H4P(dO)Rf2, and ketones,25 Rf2CdO (Rf = {CF3(CF2)5CH2CH2}) as extractants for toxic metals, radionuclides, and precious metals.26 During the course of these studies we have determined the solid state structures of several highly fluorinated molecules, and the supramolecular structures show numerous C
F 3 3 3 F
C contacts closer than the sum of the van der Waals radii (2.94 Å). This has allowed us an opportunity to examine a series of compounds that all feature the same CF3(CF2)5CH2CH2 “ponytail” by both X-ray crystallography and computational chemistry. Moreover, these weak interactions can be used to direct the assembly of selected perfluorinated compounds onto perfluorinated self-assembled monolayers (PF-SAMs).

’ EXPERIMENTAL SECTION The compounds Rf3PdO,27 Rf3PPhI,27 Rf3PdS,24 Rf2P(dO)C2H4P(dO)Rf2,24 and RfCO2H25 [Rf = CF3(CF2)nCH2CH2; n = 5, 9) were prepared according to the literature. Trioctylphosphine oxide and 1H,1H,2H,2H-perfluorodecanethiol were obtained from Aldrich and CF3(CF2)5CH2CH2I from Fluorochem and used as received. 1H, 13C, and 19F NMR spectra were recorded on a Bruker AV400 spectrometer operating at 400.13 (1H), 376.55 (19F), and 100.65 (13C) or a Bruker Avance II 600 NMR with a TCI cryoprobe spectrometer operating at 150.92 MHz (13C) and were referenced to the residual 1H or 13 C resonances of the solvent used or external CFCl3 (19F).

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F assignments were confirmed using 19F
19F COSY spectra. Mass spectra were measured on a MALDI QTOF Premier MS system. Infrared spectra of the pure compounds were recorded on a Perkin-Elmer spectrum 100 using an attenuated total internal reflection accessory. Raman spectra were obtained using 785 nm excitation on a Renishaw 1000 micro-Raman system. Melting points were measured from differential scanning calorimetry (DSC) spectra, which were recorded on a Perkin-Elmer diamond DSC. Infrared reflection absorption spectroscopy was carried out on a Bruker Tensor 27 FTIR equipped with an MCT (mercury cadmium telluride) detector using a VeeMax II specular reflectance accessory with a wire grid polarizer. All IRRAS measurements were obtained using p-polarized light at an 80
angle of incidence from the surface normal at a resolution of 4 cm
1; peak integration was carried out using WinFIRST software. Preparation of CF3(CF2)5(CH2)4(CF2)5CF3, 2. A solution of CF3(CF2)5CH2CH2I (24.5 cm3, 0.099 mol) in Et2O was added dropwise to Mg (2.43 g, 0.099 mol) over the period of 1 h. This was stirred for 2 h and cooled to 0
C. The reaction was quenched by bubbling CO2 through the solution, followed by addition of 10% H2SO4. Separation of the organic layer, drying over MgSO4, and removal of the solvent under vacuum gave a white powder that was a ca. 3:1 mixture of CF3(CF2)5CH2CH2CO2H and CF3(CF2)5CH2CH2CH2CH2(CF2)5CF3. The title compound can be isolated by repeated fractional recrystallization from CH2Cl2. Yield: 2.1 g (3.02 mmol, 3.5%). MP: 59.3
C. 1H NMR (CDCl3): δ 2.14 (t, 3JH
F = 18 Hz; CH2CF2), 1.75 (m, CH2CH2). 13C NMR (CDCl3): δ 19.9 (tt, 1JC
H = 130 Hz, 3 JC
F = 4 Hz, CH2CH2), 30.6 (tt, 1JC
H = 128 Hz, 2JC
F = 23 Hz, CH2CF2), 109.39 (tq, 1JC
F = 269 Hz, 2JC
F = 39 Hz, CF3CF2), 110.30 (tt, 1JC
F = 271 Hz, 2JC
F = 31 Hz), 111.03 (tt, 1 JC
F = 266 Hz, 2JC
F = 32 Hz), 111.07 (tt, 1JC
F = 271 Hz, 2 JC
F = 32 Hz), 117.1 (qt, 1JC
F = 287 Hz, 2JC
F = 32 Hz, CF3), 118.14 (tt, 1JC
F = 254 Hz, 2JC
F = 31 Hz). 19F (CDCl3): δ
81.36 (s, CF3),
114.93 (s, CF2CH2),
122.51 (s, CF3CF2CF2),
123.46 (s, CF3CF2CF2CF2),
124.12 (s, CH2CF2CF2),
126.74 (s, CF3CF2). IR (ATR): 2966 (w), 1709 (w), 1476 (w), 1336 (w), 1246 (m), 1218 (s), 1183 (s), 1136 (s), 1063 (s), 985 (w), 908 (w), 893 (w), 791 (w), 757 (w), 740 (w), 716 (w), 694 (s) cm
1. MS (MALDI-ToF, m/z): 695 (100%, MH+). Preparation of Perfluoro Self-Assembled Monolayers. PFSAMs of 1H,1H,2H,2H-perfluorodecanethiol were prepared according to previous reports.28 Briefly, gold substrates were immersed in 0.0010 M solutions of 1H,1H,2H,2H-perfluorodecanethiol in absolute ethanol for 1 h under argon flow. The surface was rinsed three times with absolute ethanol, sonicated for 2 min to remove any physisorbed species and dried under argon. The PF-SAM coated gold substrates were then immersed for 30 min in 0.0010 M solutions of compounds 2
7 in acetone. The substrates were rinsed three times with acetone and dried under argon prior to IRRAS characterization. Following measurements, the substrates were sonicated for 10 min in acetone to remove physisorbed species and IRRAS measurements repeated. X-ray Crystallography. X-ray diffraction measurements were made at 150 K on an Oxford Diffraction Xcalibur diffractometer equipped with a Sapphire3 detector and an Oxford Cryosystems crystal cooling device. The structure was solved and refined with the SHELXS and SHELXL suites29 using the XSEED30 interface. CCDC-831561 contains the supplementary crystallographic data for this paper. These data can be obtained free of charge from The Cambridge Crystallographic Data Centre via www. ccdc.cam.ac.uk/data_request/cif. 1436

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Chart 3. Structurally Characterized Compounds

Crystal data for 2: C16H8F26, Mr = 694.22, monoclinic, a = 13.4784(11) Å, b = 5.7846(4) Å, c = 14.5854(9) Å, β = 99.247(7)o V = 1122.40(14) Å3, Z = 2, P21/c, Dc = 2.05 g cm
3, μ = 0.272 mm
1, T = 150 K, 2578 unique reflections, 1271 with F2 > 2σ, R = 0.079, Rw = 0.15. Computational Methods. All ab initio calculations were performed using Molpro31 and employed augmented, correlation consistent basis sets aug-cc-pVnZ, where n = D, T, or Q.32 Correlation energies from ab initio methods were used to estimate the complete basis set limit by extrapolation of MP2 data from n = T and Q, following Helgaker’s scheme,33 and then applying a ΔECCSD(T)
MP2 correction using n = D.34 Density fitting, local MP2 (DF-LMP2) calculations were also employed,35 exclusively using aug-cc-pVTZ orbital and fitting basis sets. Spin component scaling for molecular interactions, SCS(MI),36 was used to correct for the known shortcomings of MP2 for noncovalent interactions dominated by dispersion forces. Local correlation and DFT data were not corrected for basis set superposition error (BSSE), following previous reports of the negligible size of such corrections with suitably large basis sets.37 All other ab initio data were corrected for BSSE using the counterpoise correction of Boys and Bernardi. DFT calculations were carried out in Turbomole38 with the def2-TZVP orbital and fitting basis sets,39 taking advantage of the resolution of identity (RI) making larger calculations viable where possible. Specific details of functionals employed are reported in relevant areas of results and discussion.40 Empirical dispersion calculations were performed using Grimme’s approach, as described in ref 41. Cartesian coordinates of relevant molecules and dimers were extracted from X-ray data, and hydrogen positions were adjusted to standard lengths from neutron diffraction data along the X—H vector.

’ RESULTS AND DISCUSSION To give a comprehensive description of the interactions involved in the fluorinated compounds we have available, we first discuss the supramolecular structures of three compounds characterized by X-ray diffraction. We then use computational chemistry to gain a theoretical insight into these interactions. Two areas will be explored in detail, first the stabilization energies of the interactions, followed by an investigation into the bonding characteristics utilizing an AIM approach. Finally, we show that these interactions can be used to direct the assembly of a series of compounds onto perfluorinated self-assembled monolayers. Synthesis and Solid State Structures. We have previously reported on the preparation and solid state structures of two compounds (125 and 326) that feature a fluorous ponytail but did not discuss in detail the supramolecular structure of these. Herein, we present the crystal structure of 2 and describe the differences in packing between these three compounds (Chart 3). 1 exhibits the familiar head-to-head hydrogen bonding between carboxylic acid groups,42 as well as numerous C
F 3 3 3 F
C and C
F 3 3 3 H
C contacts (Figure 1). The metric parameters

Figure 1. Packing Diagram of 1 reproduced with permission from Reference 25.

of the C
F 3 3 3 H
C interaction are dF 3 3 3 H = 2.624 Å and — C
F 3 3 3 H—C = 143.1
and 106.24
. Substantially lower torsion angles have been previously noted in fluorous chains compared to hydrocarbon chains to relieve electrostatic interactions between two fluorine groups, although other factors may also be involved43 (average C
C—C—C = 166
), which imparts a slight twist to the fluorocarbon chain. One helix then packs against the opposite orientation in an adjacent fluorocarbon chain, consistent with previous reports.44 There are also close contacts between two fluorous chains that involve a three point interaction between one fluorine atom on one molecule and two fluorine atoms on a second. The metric parameters are dF 3 3 3 F = 2.89 Å and — F 3 3 3 F 3 3 3 F = 62
. A CCDC search reveals that these interactions are not common (53 hits with the following restraints: dF 3 3 3 F less than the van der Waals radii; — F 3 3 3 F 3 3 3 F = 40
90
; Figure S1, Supporting Information) and have not been commented on. To determine if these are specific interactions or just a consequence of close-packing, and to gain further insight into the degree of stabilization, it was therefore probed by computational chemistry (vide infra). The F-alkyl/alkyl triblock45 molecule 2 can be synthesized from a Wurtz coupling of the Grignard reagent CF3(CF2)5CH2CH2MgI. We have isolated this by repeated fractional crystallization from a mixture of the corresponding carboxylic acid CF3(CF2)5CH2CH2CO2H and 2. This has not been fully characterized by NMR spectroscopy before, so we have included complete NMR data in the Experimental Section. The solid state structure of 2 has not been previously reported, although it has been characterized by powder diffraction46 and thermal techniques.47 In line with the powder diffraction study we have observed two different polymorphs;48 only one gave a solution, which corresponds to the β-polymorph, and is shown in Figure 2a. The bond lengths and angles are collated in the Supporting Information (Table S1). The bond angles are as would be expected (C
C
C = 111.8
116.8
, F
C
F = 106.2
109.4
), and the C
C bond lengths in the fluorous domain (1.533
1.560 Å) are slightly longer than in the hydrocarbon domain (1.515(8) and 1.523(5) Å). The packing diagram of 2 is shown in Figure 2b. Interestingly, there are no C
H 3 3 3 F
C interactions in the supramolecular structure, presumably as the separation between chains are determined by the longer C
F 3 3 3 F
C contacts, by virtue of fluorine having a larger van der Waals radius. There are type I C
F 3 3 3 F
C close contacts from the
CF3 groups that form infinite linear chains in the b-direction of the crystal cell. The metric parameters for these interactions are dF 3 3 3 F = 2.897 Å and — C
F 3 3 3 F
C = 141.9
. The most obvious feature of the solid state structure of 3 is the hydrogen bonding between the C
H protons α to 1437

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Figure 2. Ortep plot (thermal ellipsoids at 50% probability) (a) and packing diagram (b) of 2.

Figure 3. Packing Diagram of 3. Reproduced from ref 26.

the phosphorus and the phosphine oxide (dC 3 3 3 O = 3.38 Å, — C
H 3 3 3 O = 145
), which form infinite stacks along the a-axis of the crystal cell (Figure 3). This motif is common in phosphine oxides structurally characterized, with a CCDC search revealing C 3 3 3 O distances in the range 3.046
3.420 Å. There are a number of type I C
F 3 3 3 F
C and C
H 3 3 3 F
C (dF 3 3 3 H = 2.609 Å, — C
F 3 3 3 H—C = 147.59
and 148.17
) interactions that form fluorous and non-fluorous domains. Interestingly, as with 2, there are close contacts between two fluorous chains that involve a three point interaction between one fluorine atom on one molecule and two fluorine atoms on a second ponytail. The metric parameters are dF 3 3 3 F = 2.87 Å and — F 3 3 3 F 3 3 3 F = 66
. To compare the degree of stabilization with the same interactions found in 1, these were studied by computational chemistry (vide infra). From the three crystal structures studied, it is clear that there are a number of contacts shorter than the van der Waals radii. In addition to C
H 3 3 3 O and C
H 3 3 3 F
C interactions, there are also C
F 3 3 3 F
C interactions that could be classed as type I or type II, and a new class of interaction between two ponytails. Computational chemistry has been extensively used to examine these weak interactions and can give information on the degree of stabilization and, via the AIM approach, a more thorough

Figure 4. CF4 dimer in D3d conformation.

understanding of nature of these interactions. As the molecules are computationally large, we have used the X-ray structures as a starting point for the computations. Computational Studies. As a small, computationally tractable model of C
F 3 3 3 F
C interactions, we followed the work of Mahlanen et al. and employed the dimer of CF4.49 Previous work shows that the D3d form of this dimer with staggered
CF3 groups in close contact, as shown in Figure 4, is the most stable. We initially performed a scan of intermolecular separation using DF-LMP2 methods, as shown in Supporting Information (Figure S2), which shows stabilization of the dimer at separations larger than 3.6 Å, reaching an energy minimum at a C 3 3 3 C 1438

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Table 1. Interaction Energy of (CF4)2 at SCS(MI) Minimum Energy Geometry (kcal/mol) method

a

Table 2. Interaction Energies of Dimers Taken from Crystal Structures (kcal/mol)

ΔE

B97-Da

CBS(T)


0.77

1a


13.23 (
3.19)

MP2/CBS MP2/aug-cc-pVTZ


0.69
0.60

1b 1c


1.69 (
5.86)
0.58 (
0.85)

DF-LMP2a


0.63

2a


4.66 (
6.21)

SCS(MI)-DF-LMP2a


0.48

2b


1.31 (
1.77)

B97-Da


0.61

2c


0.26 (
0.60)

B2-PLYPa


0.20

3a


29.64 (
20.82)

BHandH


0.74

3b


1.82 (
2.65)

B3LYP

+0.45

3c


6.55 (
9.89)

HF

+0.49

3d


5.76 (
10.00)

Not counterpoise corrected.

separation of 4.1 Å. This geometry was then used for all ab initio calculations of interaction energy, the results of which are reported in Table 1. Our best estimate of the interaction energy in this model system is the CBS(T) value of
0.77 kcal/mol, obtained from extrapolation of MP2 energies to the basis set limit plus correction from CCSD(T) calculation. This stabilization is similar to the value quoted by Mahlanen et al. and comparable in size to other weak noncovalent interactions such as C
H 3 3 3 O and C
H 3 3 3 π hydrogen bonds. For example, C
H 3 3 3 H
C interactions in methane were recently reported to also have the most stable orientation as the D3d staggered form, with interaction energy of
0.42 kcal/mol at MP2/6-311++G(3df,3pd) level.50 Methods based on MP2 are in good general agreement with the CBS(T) value, although SCS(MI) scaling of MP2 data actually gives worse agreement than unscaled data. DFT methods such as B97-D, which includes an empirical correction for dispersion energy, or the “double-hybrid” B2-PLYP, are also close to the best estimate. The excellent performance of Becke’s half-and-half (BHandH) functional, previously noted to work well for stacking interactions, is slightly surprising, and may well be fortuitous. In contrast, the popular B3LYP functional fails to give any stabilization of this dimer, as does the uncorrelated HF method. A feature of the local correlation approach is that it allows decomposition of intermolecular correlation energy into contributions from dispersion and ionic excitations, thereby giving deeper physical insight into the origin of noncovalent interactions. In this case, partition of DF-LMP2 data indicates that dispersion forces dominate the interaction, resulting in
0.95 kcal/mol of stabilization, in comparison to ionic interactions that yield just
0.11 kcal/mol. The total intermolecular correlation energy in the DF-LMP2 approach (
1.06 kcal/mol) compares reasonably well with the empirical dispersion correction of
1.77 kcal/mol obtained from the B97-D approach, confirming the suitability of the latter for description of larger systems. In LMP2, total ΔE ≈ dispersion + ionic + HF: as a consistency check, this gives
0.63 ≈ (
0.95
0.11 + 0.49) =
0.57 kcal/mol, so remaining terms are essentially negligible. A series of dimers was extracted from the crystal structures of 1, 2, and 3 by applying individual symmetry operations to the entire molecule. The size of these dimers (over 150 atoms and 4000 basis functions) means that only B97-D data could be calculated, with binding energies reported in Table 2. Using this methodology, three dimers were extracted from the crystal structure of 1 (Figure 5). The first, 1a, contains two hydrogen

a

Contribution from empirical dispersion correction in parentheses.

Figure 5. Dimers taken from crystal structures of 1 and 2.

bonds linking carboxylic acid groups and no close C
F 3 3 3 F
C contacts. 1b exhibits three C
F 3 3 3 F
C contacts of length 2.850, 2.903, and 2.890 Å as well as one C
F 3 3 3 H
C contact of length 2.629 Å, whereas 1c contains just one C
F 3 3 3 F
C contact at 2.911 Å. As expected, the hydrogen bonded dimer 1a is very strongly stabilized with only a minor contribution from dispersion forces. Although this dimer is not of direct relevance to the main focus of this study, it provides a useful reference point for comparison. 1b is only weakly stabilized despite its numerous close contacts, but these data confirm that the three-point interaction motif discussed above is stable. Binding of 1c is very weak indeed with a binding energy close to that of (CF4)2. 2 forms numerous C
F 3 3 3 F
C contacts in the solid state, from which a further three dimers were extracted, as shown in Figure 5. Two such dimers, 2a and 2b, are generated by 2-fold screw axes to give approximately perpendicular arrangements, whereas the third stems from translation giving an end-to-end orientation, 2c. 2a contains two C
F 3 3 3 F
C contacts of length 2.864 and 3.128 Å, whereas 2b and 2c exhibit just one such interaction each, with lengths of 2.895 and 2.897 Å, respectively. Interaction energies at B97-D level are reported in Table 3. The stabilization of 2a is again dominated by dispersion forces, whereas 2b and 2c are much more weakly bound. 2c features the type I interaction, which in the case of uracil derivatives gave a stabilization energy of ca. 1 kcal/mol.16 1439

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Table 3. Summary of AIM Analysisa F3 3 3F

C
H 3 3 3 F C
H 3 3 3 O O
H 3 3 3 O O 3 3 3 O 2, 0.048

1a 1b 3, 0.007

O3 3 3F

2, 0.009

1, 0.003

1, 0.012 1, 0.005

1c 1, 0.005 2a 4, 0.014

7, 0.018

2b 2, 0.008 2c 1, 0.005 3a 23, 0.082

3, 0.015

3b 3, 0.018 3c 11, 0.044

5, 0.016

3d 11, 0.057

4, 0.016

3, 0.028

a

Reported as number of bond critical points located and sum of electron density at each (e 3 bohr
3).

Figure 7. Intermolecular bond paths located in dimer 1b, showing three-point interaction (right). Electron density at selected BCP’s (e 3 bohr
3) is also shown.

Figure 6. Dimers taken from crystal structure of 3.

3a is generated by translation along the a-axis of the crystal cell and contains three PdO 3 3 3 H
C contacts, as well as numerous C
F 3 3 3 F
C and C
H 3 3 3 F
C interactions (Figure 6). This dimer exhibits very strong stabilization, largely but not entirely due to dispersion interactions. For comparison, a dimer of Et3PdO, constructed from the same geometry by truncating fluorous chains to CH3, is bound by
7.83 kcal/mol (
5.25 kcal/mol from dispersion), suggesting that the fluorous chains contribute more than 20 kcal/mol of stabilization of 3a. This is probed in more detail below, but indirect experimental evidence supporting this stabilization comes from a comparison of the melting points of 3 (64.41
C) and the nonfluorous analogue trioctylphosphine oxide (54.83
C) as determined by DSC.45 3b contains a terminal
CF3 group interacting with
CF3 and
CF2 groups in a second molecule, in a manner similar to that for (CF4)2, and is only very weakly stabilized. 3c and 3d contain numerous C
F 3 3 3 F
C contacts and are stabilized by ca. 6 kcal/mol in total, essentially all of which stems from dispersion forces. The energy calculations described above include all interactions in the dimer and do not address specific interactions. These can be probed in more detail with atoms-in-molecules (AIM) analysis, in which unambiguous intermolecular contacts can be

identified through the presence of (3,
1) or bond critical points (BCPs). Properties at such BCPs are widely used to characterize covalent and noncovalent interactions, for instance via correlations between the electron density at the BCP and the strength of interaction. Table 3 summarizes this analysis and shows great diversity of intermolecular interactions in the dimers considered. Dimer 3a exhibits over twenty C
F 3 3 3 F
C interactions, along with three each of C
H 3 3 3 F
C and C
H 3 3 3 O paths, whereas 3c and 3d each exhibit eleven C
F 3 3 3 F
C interactions along with several C
H 3 3 3 F
C ones. In contrast, 3b shows just three C
F 3 3 3 F
C contacts, 2b two such contacts, and the end-toend dimers 1c and 2c a single F 3 3 3 F interaction each. The most complex topology is found for 1b, which contains not only C
F 3 3 3 F
C, C
H 3 3 3 F
C, and C
H 3 3 3 O interactions but also O 3 3 3 O and O 3 3 3 F contacts (Figure 7). Table 3 and Figure 7 also report the electron density BCPs: the density in conventional hydrogen bonds (1a) is much greater than in C
F 3 3 3 F
C, C
H 3 3 3 F
C, or C
H 3 3 3 O ones. However, the large number of weak interactions leads to collectively large amounts of electron density at intermolecular BCPs, most notably in 3a but also in 3c, 3d, and 2a. In dimer 3a in particular, the cumulative effect of numerous weak interactions leads to the greater stabilization of this dimer over the conventionally hydrogen-bonded 1a, or the model dimer of Et3PdO. It is worth noting that the Laplacian (r2) at the BCP is positive in all interactions studies, as expected for weak interactions. The electron density at the BCPs compare well to those reported in the literature, for example, H
F 3 3 3 F
F,51 C
F 3 3 3 F
C in difluorinated aromatics,52 or H3C
H 3 3 3 F
CH3.53 Finally, there does appear to be a correlation between the electron density at BCP and the calculated stabilization energy (Figure S4, Supporting Information). In summary, the computational studies reveal that the plethora of weak interactions present in these systems can be safely characterized as weak, but stabilizing, interactions. Hydrogen Bonding in 3. As the hydrogen bonding in 3 shows substantial stabilization energy, we have examined this in more detail, both computationally and experimentally. The calculated molecular dipole moment of 3 is 4.89 D, only slightly greater than that of Et3PdO (4.45 D). Atomic charges calculated via the 1440

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Figure 8. Infrared and Raman spectra of the C
H stretching region of Oct3PdO (left), 3 (middle), and Rf3PdS (right).

natural bond orbital (NBO) scheme indicates that the nature of the PdO bond is barely perturbed by the presence of fluorous chains (Table S3, Supporting Information). Similarly, the three C atoms α to the phosphorus carry very similar charges in both species. A more significant difference is apparent in the charges on hydrogens bound to these carbons, which are uniformly more positive in 3 than in the nonfluorinated analogue. This change is also evident in plots of the molecular electrostatic potential projected onto an electron density isosurface (Figure S5, Supporting Information). Both plots show similar negative potentials associated with the PdO group, but in 3 regions of positive potential are associated with CH2 groups that are not present in Et3PdO. This in turn suggests that the strength of the C
H 3 3 3 O interactions in 3a stems from polarization of the C
H bonds involved, rather than of the PdO bond. It should be noted that experimental studies on metal complexes of the phosphines {CF3(CF2)n(CH2)m}3P have shown that an ethyl insulating group does moderate the electron withdrawing effect of the perfluoro-chain to a high degree.54 Commonly the degree of hydrogen bonding is explored spectroscopically using infrared spectroscopy. Although the C
H stretches in the infrared spectra are very weak in fluorinated compounds as the polarity of the C
H bond increases,55 the corresponding bands in the Raman spectra are not as sensitive to this effect.56 Additionally, bands associated with Fermi resonances and overtone bands are not observed in fluorinated chains due to the shift in frequencies of the CH2 bending mode because of inductive effects from the CF2 groups, so the number of peaks is lower than in hydrocarbon chains.57 Figure 8 shows the IR and Raman spectrum of the C
H stretching region of 3 and, for comparison, Oct3PdO (trioctylphosphine oxide) and Rf3PdS (Rf = CF3(CF2)5CH2CH2) to deconvolute the effects of the electron withdrawing perfluorinated chain and the degree of hydrogen bonding. The analogous phosphine sulfide has been included as little hydrogen bonding

would be expected as the PdS function would not be a good hydrogen bond acceptor.58 The bands in the PdS spectrum can be readily assigned to the fundamental symmetric and antisymmetric C
H stretch. In the spectra of 3 it can be seen that there are more bands and a blue shift of 17 cm
1 in the Raman and 23 cm
1 in the IR spectra, which suggests that hydrogen bonding is present. A comparable blue shift in the IR and Raman spectra of the hydrogen bonded complex of Me2O and CHF3 was reported.59 The spectrum of Oct3PdO shows the fundamental C
H stretching modes for both the CH2 and CH3 groups, which are slightly blue-shifted in comparison to 3, which may be due to either the electron withdrawing perfluoro-chain or to C
H 3 3 3 O hydrogen bonding. DFT calculations on 3 and Et3PdO supports the former cause, revealing a blue shift of 15 cm
1 upon fluorination. In addition to examining the C
H stretch we can clearly observe the PdO stretch in the Raman spectra. PdO stretching bands in the experimental spectra are observed at 1155 cm
1 for 3 and 1144 cm
1 for Oct3PdO, with analogous DFT values of 1154 cm
1 for 3 and 1148 cm
1 for Et3PdO. These values corroborate the idea that the PdO bond is barely affected by the presence of fluorinated chains, and instead that it is polarization of the C
H bonds that is responsible for the stabilization. Physisorption Studies. Computational results and solid state structures demonstrate that C
F 3 3 3 F
C interactions are energetically favorable; therefore, we investigated their applicability to solid phase extraction via physisorption onto fluorinated surfaces. For this purpose, we investigated the adsorption of 2
7 (Charts 3 and 4) onto a model fluorinated surface prepared via self-assembly of 1H,1H,2H,2H-perfluorodecanethiol (PFSAM) on a gold surface, using infrared reflection absorption spectroscopy (IRRAS). These compounds were chosen to investigate the effect of chain length, the number of ponytails, and the charge on the molecules; thus 3 and 4 show only an increase in 1441

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Chart 4. Perfluorinated Compounds Physisorbed onto Perfluorinated Self Assembled Monolayers (PF-SAM)

Figure 10. Integrated absorbances in the region 1400
1000 cm
1 of compounds 2
7. Error bars indicate 90% C.I.

Figure 9. Top: IR absorption spectrum of 7. Bottom: IRRAS spectra of PF-SAM (A), of PF-SAM after adsorption of compound 7 (B), and after vigorous sonication to remove physisorbed material (C). Spectra have been offset for clarity.

the number of
CF2 moieties, whereas 7, 2, 3, and 6 show an increase in the number of ponytails. All of the fluorinated compounds discussed in this work display strong infrared absorption peaks in the region 1000
1400 cm
1, which are assigned to C
F stretching modes of the fluorinated chains.60 As an example, Figure 9 (top) shows a typical absorption infrared spectrum in this region for compound 7. The spectrum displays several unresolved peaks with two prominent absorptions at 1148 and 1200 cm
1 that, based on previous assignments,61 are attributed to the symmetric and antisymmetric
CF2 stretching modes, respectively, of the fluorinated chains. Three peaks with maxima at 1342, 1357, and 1374 cm
1 are attributed to
CF3 stretching modes. Peak intensities observed in IRRAS spectra are the result of both surface coverage and orientation effects; however, in the case of disordered films, the integrated intensity can be used to monitor surface coverage.62 We therefore used the integrated peak area to quantitate the adsorption of fluorinated compounds on the PF-SAMs on gold. Figure 9 (bottom) shows a typical set of IRRAS spectra showing absorptions in the 1000
1400 cm
1 region for the PF-SAM (trace A), the surface after exposure to acetone solutions of 7 (trace B) and the same surface after 10 min

sonication in acetone to remove physisorbed material (trace C). IRRAS spectra of PF-SAMs display absorptions in the 1400
1000 cm
1 with an integrated area of 0.18 ( 0.01 (90% C.I.). Upon adsorption of fluorinated compounds, the integrated intensity in the region 1000
1400 cm
1 increases as expected from an increase in the surface density of
CF2 and
CF3 groups. Control experiments on bare gold surfaces did not yield significant absorption peaks in this region, thus indicating that an increase in infrared absorption results from molecular assembly onto the PFSAM, as opposed to simple precipitation/aggregation from solution. Following sonication, the intensity of the C
F stretching absorbances returns to the initial values, suggesting that physisorption on fluorinated surfaces is reversible. Figure 10 shows a summary of the integrated absorptions measured for compounds 2
7: it was consistently observed that all of the compounds physisorbed on the PF-SAM. The largest increase in surface coverage is observed for the bidentate phosphine 6, which has four fluorinated
(CF2)5CF3 chains. The increase in adsorption is much higher than that expected on the basis of the results for 3, which possesses three of the same fluorinated chains (expected approximately 30% increase). This would suggest that either the arrangement of the fluorinated chains or the presence of PdO groups is responsible for more favorable interactions with the surface. Control experiments carried out using Oct3PdO, the nonfluorous analogue of 3, showed no increase in intensity in the C
H (2800
3000 cm
1) stretching region, thus suggesting that PdO groups should not promote adsorption significantly. Therefore, we propose that the arrangement of fluorinated chains is important in determining adsorption properties. It is interesting to note that 6 would probably exist in the transoid form, as seen crystallographically for other examples of bidentate phosphine oxides,63 which may give a bigger “footprint” compared to 3. 4 and 5 both have the same number of fluorous groups so the increase in 5 compared to 4 may be due to either electrostatics or the phenyl groups allowing for a further stabilization by π
π interactions to form a bilayer rather than monolayer motif. Compounds 3, 2, and 7 show a steady decrease in
CF2 and
CF3 functionality that is 1442

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The Journal of Physical Chemistry A mirrored in decreased absorbance observed for their reactions on the fluorinated thiol coated surface.

’ SUMMARY In conclusion we have characterized C
F 3 3 3 F
C interactions in the solid state via X-ray crystallography and comprehensive computational studies. We have found a previously unreported three point type of C
F 3 3 3 F
C interaction. The interaction energies are from 1 to 20 kcal/mol, which are strong enough to physisorb selected fluorous molecules onto fluorous self-assembled monolayers; this has been qualified by IRRAS techniques. This suggests that it is possible to design solid phase recovery systems for ligands, catalysts, or pollutants based on fluorous interactions. The weakness of fluorous interactions relative to other intermolecular forces is an advantage for such applications because, as shown in our experiments, it allows for ease of desorption/recovery. This work adds to the growing body of evidence for these weak interactions and should be considered as another addition to the supramolecular chemists’ toolbox. ’ ASSOCIATED CONTENT

bS

Supporting Information. Full crystallographic descriptions (CIF), atomic data, and computational details. This material is available free of charge via the Internet at http://pubs.acs. org.

’ AUTHOR INFORMATION Corresponding Author

*Tel: +353-1-8963501. Fax: +353-1-6712826. E-mail: bakerrj@ tcd.ie.

’ ACKNOWLEDGMENT We thank TCD and the Leverhulme Trust for funding, and Cardiff’s Advanced Research Computing facility ARCCA for computing resources. Part of the work in this publication has emanated from research conducted with the financial support of Science Foundation Ireland under Grant Number 09/RFP/ CAP2174. ’ REFERENCES (1) Gilli, G.; Gilli, P. The Nature of the Hydrogen Bond; Oxford University Press: Oxford U.K. 2009. (2) (a) Hunter, C. A.; Sanders, J. K. M. J. Am. Chem. Soc. 1990, 112, 5525–5534. (b) Burley, S. K.; Petsko, G. A. Science 1985, 229, 23–28. (c) Kim, K. S.; Tarakeshwar, P.; Lee, J. Y. Chem. Rev. 2000, 100, 4145– 4186. (d) Lee, E. C.; Kim, D.; Jurecka, P.; Tarakeshwar, P.; Hobza, P.; Kim, K. S. J. Phys. Chem. A 2007, 111, 3446–3457. (e) Singh, N. J.; Min, S. K.; Kim, D. Y.; Kim, K. S. J. Chem. Theory Comput. 2009, 5, 515–529. (f) Janiak, C. J. J. Chem. Soc., Dalton Trans. 2000, 3885–3896. (g) Meyer, E. A.; Castellano, R. K.; Diederich, F. Angew. Chem., Int. Ed. 2003, 42, 1210–1250. (3) (a) Schottel, B. L.; Chifotides, H. T.; Dunbar, K. R. Chem. Soc. Rev. 2008, 37, 68–83. (b) Frontera, A.; Gamez, P.; Mascal, M.; Mooibroek, T. J.; Reedijk, J. Angew. Chem., Int. Ed. 2011, 50, 9564–9583. (4) (a) Ma, J. C.; Dougherty, D. A. Chem. Rev. 1997, 97, 1303–1324. (b) Dougherty, D. A. Science 1996, 271, 163–168. (5) Mooibroek, T. J.; Gamez, P.; Reedijk, J. CrystEngComm 2008, 10, 1501–1515. (6) Nishio, M.; Hirota, M.; Umezawa, Y. The CH-π Interaction: Evidence, Nature and Consequences; Wiley-VCH: New York, 1998.

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