Water–Halogen Interactions in Chlorine and ... - ACS Publications

Article pubs.acs.org/JPCC

Water−Halogen Interactions in Chlorine and Bromine Clathrate Hydrates: An Example of Multidirectional Halogen Bonding Konstantin A. Udachin,† Saman Alavi,†,‡ and John A. Ripmeester*,† †

National Research Council Canada, Ottawa, Ontario K1A 0R6 Canada Department of Chemistry, University of Ottawa, Ottawa, Ontario K1N 6N5, Canada

‡

S Supporting Information *

ABSTRACT: There are unresolved questions in our understanding of the clathrate hydrates with halogen molecule guests. In this work, we describe the synthesis and single crystal X-ray structural determination of cubic structure I (sI) Cl2, tetragonal structure III (sIII) Br2, and mixed cubic sI Cl2− Br2 clathrate hydrate phases. Unlike the previous X-ray structural determination of the pure chlorine clathrate hydrate by Pauling and Marsh (Proc. Nat. Acad. Sci. USA 1952, 38, 112), we report that 32.5% of the clathrate small cages are occupied by Cl2 molecules. Pure Br2 produces the unique tetragonal sIII clathrate hydrate whose structure is refined from single-crystal measurements, identifying disordered Br2 guest positions in cages and giving structural details not previously identified. In mixed sI clathrate hydrate of Cl2 and Br2, we observed that Br2 guests fit into the large sI cages. The introduction of Br2 in the Cl2 hydrate lattice causes severe distortions because of strong Br···OH2 interactions. In these three hydrate structures, based on accepted van der Waals radii, the effective diameters of the clathrate hydrate cages are smaller than the size of the Cl2 and Br2 guest molecules. The accommodation of these guests in the cages and the short O···Cl and O···Br contact distances reveal halogen bonding between the water oxygen atoms and the dihalogen guests. The X-ray structural determinations of the length of the Cl−Cl and Br−Br bonds of the Cl2 and Br2 guests show significant librational shortening compared to free Cl2 and Br2 molecules.

■

INTRODUCTION Clathrate hydrates are guest−host compounds where the hydrogen bonded three-dimensional water lattice forms cages which encapsulate guest species. Three clathrate hydrate structures are of interest in this work: cubic structures I, II (sI, sII), and the tetragonal structure III.1,2 Almost all atomic or molecular species that can fit into hydrate cages can be clathrate hydrate guest molecules, so there are relatively few restrictions other than size.3 Classically, clathrate hydrate stability was considered to depend on short-range guest−host repulsive (steric) interactions from hydrophobic species,4 but clathrate hydrates of many water-miscible guests, including organics,5 bases, and mineral acids,6 are known as well. This indicates that for neutral guest species, short-range attractive interactions (such as hydrogen bonding) between the guests and water can be present in the clathrate hydrates, in addition to steric repulsions. Indirect evidence for these strong hydrogen bonds was known for many years.4,7−9 Only recently have advances in structural investigations by single-crystal X-ray crystallography and molecular dynamics methods shed new light on the complex disorder and guest− host interactions in clathrate hydrates phases.9−11 Guest−host hydrogen bonds of the type guest-O···H−OH and guestN···H−OH, ranging from transient to long-lived, were identified by molecular dynamics (MD) simulations, single© 2013 American Chemical Society

crystal X-ray diffraction, IR and NMR spectroscopic measurements. Dynamic hydrogen bonds with different formation probabilities and lifetimes and even cases where a guest molecule displaces water from the normal position in the hydrate lattice were observed.12 An interesting feature of these materials is that hydrogen-bonding, taken normally to be a directional, specific interaction capable of determining structural motifs in crystal engineering, becomes a nondirectional, or perhaps better stated, a multidirectional guest−host interaction because of the complex disorder and guest dynamics in the clathrate hydrate cages. The dihalogen molecules, chlorine and bromine, are interesting guests that form clathrate hydrates for which some structural information is available. Indeed, the first clathrate hydrate discovered, that of chlorine, was reported about 200 years ago in 1811 by Humphrey Davy13 and its halogen-towater stoichiometry studied by Michael Faraday 12 years later.14 This hydrate forms when bubbling chlorine gas into cold water. The clathrate nature of chlorine hydrate was not revealed until 1953 when three groups independently reported crystal structures of sI and sII hydrates. Pauling and Marsh15 Received: March 8, 2013 Revised: June 10, 2013 Published: June 11, 2013 14176

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

reported the structure of chlorine hydrate as cubic sI which has a unit cell designated as 6T·2D·46H2O where the T is a tetrakaidecahedral large cage (51262) and D is a dodecahedral small cage (512). On the basis of cage sizes and accepted van der Waals radii, Pauling and Marsh assumed that the Cl2 guests should reside only in the large cages, giving each chlorine molecule a hydration number of 46/6 =7.67 in the clathrate hydrate. Based on the Cl−Cl bond length of 1.99 Å and the van der Waals radius of 1.75 Å for each chlorine atom, the van der Waals length of Cl2 is 5.49 Å. The free diameter of the sI large cage is usually taken to be 5.8 Å, so the guest−host fit is reasonable. It would appear that Cl2 is far too large to fit into the small D cages, which are usually given a free diameter of 5.0 Å.2 However, the above structural assignment did not match experimental hydration numbers of Cl 2 ·7.27H 2 O and Cl2·6.04H2O for the clathrate hydrates as reported by Glew and Rath16 and Anwar-Ullah17 which require 16% and 81% small cage filling with Cl2, respectively. In 1983 Cady18 performed careful gravimetric studies of Cl2 uptake into the hydrate phase at different gas pressures and determined that the stoichiometric ratio of Cl2-to-water is consistent with partial filling of the small sI cages. Cady determined that the experimental gas uptake as a function of external Cl2 pressure, which is related to the degree of cage filling by the Cl2 guests, was larger than predicted by the van der Waals−Platteeuw theory. This larger gas uptake may be due to the halogen bonding between Cl2 and water, which enhances the strength of interactions between these molecules (see below). Bromine hydrate also was one of the first hydrates reported.19 Although bromine hydrate was originally assumed to be sI,20 Jeffrey suggested a unique tetragonal structure, which he classified as structure III (sIII).21 The unit cell of sIII is designated as 4P·16T·10D·172H2O, where the P is the large pentakaidecahedral cage with 12 pentagonal and 3 hexagonal faces (51263). On the basis of classical thermodynamic measurements, crystal color, and morphology, Dyadin and coworkers22 suggested up to four different structures for this hydrate. However, single-crystal X-ray diffraction methods23 showed that the crystals of different morphology and color were all consistent with Jeffrey’s original suggested tetragonal structure, the hydration number (i.e., Br2 cage occupancy) being the chief variable. Other recent work has again presented the possibility of structures in addition to sIII.24,25 To our knowledge, fluorine and iodine hydrates have not been synthesized, although an sI hydrate of BrCl has been reported.19,20 The interactions of dihalogen molecules with water and other organic molecules were studied from the 1950s by Hassel and co-workers. They interpreted observed short distances between the dihalogen molecules and organic molecules in complexes and cocrystals as signs of interatomic charge transfer interactions.26 Some Br···C, Br···O, Cl···C, and Cl···O distances in cocrystals determined by Hassel and co-workers (given in Table 1) are shorter than the sum of van der Waals radii of the halogen atom with the C or O atoms of the organic molecule, but longer than the sum of the corresponding covalent radii of X−C and X−O bonds. In these cocrystals, the X−X bond lengths are longer than those of the free halogen molecules. Spectroscopic and computational studies indicate strong dihalogen-water interactions. Spectroscopic studies show that the 79Br−81Br stretch frequency red shifts by 17 cm−1 in going

Table 1. Experimentally Determined Dihalogen−Organic Molecule Distances26 complex

distance/Å

benzene - Br2 p-dioxane - Br2 benzene - Cl2 p-dioxane - Cl2 acetone - Br2

2.28 2.31; 2.7 1.99 2.02 2.28

from the gas phase to the liquid aqueous phase and by 2 cm−1 upon incorporation into the clathrate hydrate phase.27,28 Janda and co-workers measured the UV−visible spectrum of the pure Br2 clathrate hydrate.24 In addition to electronic transitions of the bromine molecule, which are similar to those of gas phase, Br2 in the clathrate hydrate phase shows a signature high-intensity charge transfer band. High-level computational studies of X−X vibrational frequencies in linear gas phase X2···OH2 complexes of Cl2 and Br2 predict red shifts of 17 and 10 cm−1 for vibrational stretch bands of Cl2 and Br2, respectively.28 This red shift increases up to 60 cm−1 in complexes with up to five water molecules. The red shifts are accompanied by elongation of the Cl−Cl and Br−Br bonds and have been attributed to electron donation from the water oxygen lone pair site (Lewis base) to the X2 antibonding σ* orbitals of the dihalogen molecule (Lewis acid). This is an example of the well-known halogen bonding.29−31 Several quantum chemical studies of halogen-water complexes, X2···(H2O)n, have been performed.28,31 In these complexes, one halogen atom initially interacts with the oxygen lone pair electrons in water. Additional water molecules preferentially interact with this first water molecule and not the second halogen atom. In larger water−halogen clusters, asymmetric water interactions between different water molecules with the two halogen atoms are observed. Some X−X and X−X···OH2 bond lengths for the chlorine− and bromine−water complexes are given in Tables 2 and 3. The X···O distances in these complexes range from 2.6 to 2.8 Å for chlorine and 2.75 to 2.85 Å for bromine. These distances are considerably shorter than the sum of van der Waals radii of the halogen and water oxygen atoms, which is 3.04 Å for Cl···O and 3.37 Å for Br···O interactions. The binding energies of the halogens with the water clusters, Ebind = E(X2·H2O) − E[H2O] − E(X2), range between 2.8 for Cl2 and 3.9 kcal·mol−1 for Br2 complexes. Schofield and Jordan performed molecular dynamics simulations on the Br2 tetragonal sIII and cubic sI and sII phase clathrate hydrate.25 To model the complex charge distribution in Br2, they developed a intermolecular potential with electrostatic charges, dipole moments, and quadrupole moments on the Br atoms and a charge on the molecule center of mass. The center of mass was also assigned a charge on a spring (COS)-type term to model the molecule polarizability. The bromine−water and bromine−bromine interactions were also assigned exchange repulsion, dispersion, and charge penetration terms. The H2O molecules in the clathrate hydrate phase were also assigned COS potential terms. Using this potential to calculate the energies of different bromine clathrate hydrate phases, Schofield and Jordan concluded that the tetragonal sIII gives the lowest energy Br2 clathrate hydrate, but found that the energies of the cubic sII hydrate could also be accessible to the bromine guests under proper conditions. 14177

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

interactions is a noncovalent, structure-directing interaction, distinct from hydrogen bonding. Halogen bonding was first noted many years ago26 and has attracted attention in crystal engineering in the past 10 years.29,30 In halogen bonding, the halogen atom has a σ-hole along the A−X bond, pointing away from the A atom. This σ-hole functions as the site for accepting electron density from the D atom. Some of the charge density is ultimately transferred to the A atom.

Table 2. Chlorine Bond Distances (Å) for the Free Molecule, Cl−Cl···OH2 Complex, and in the Clathrate Hydrate Phasesa distance

d(H2O···X−X)

d(Cl−Cl) (free) d(Cl−Cl) (H2O···Cl−Cl) d(Cl−Cl) (sI D cage, expt.) d(Cl−Cl) (sI T cage, expt.) d(Cl−Cl) (sI mixed Cl2/Br2 D cage, expt.) d(Cl−Cl) (sI mixed Cl2/Br2 T cage, expt.) d(H2O···Cl−Cl) (gas complex, calc.) d(H2O···Cl−Cl) (sI D cage, expt.) d(H2O···Cl−Cl) (sI T cage, expt.) d(H2O···Cl−Cl) (sI Cl2/Br2 D cage, expt.) d(H2O···Cl−Cl) (sI Cl2/Br2 T cage, expt.)

2.00b 2.050,c 2.031d 1.870(16) 1.883(9), 1.888(10) 1.865(18) 1.861(9) 2.755,c 2.799d 2.902−2.97 3.073−3.20 2.77, 2.84 2.70, 2.72

■

RESULTS AND DISCUSSION Single-Crystal X-ray Structural Determination of Clathrate Hydrate Phases. Chlorine Clathrate Hydrate. As shown previously by Pauling and Marsh, chlorine forms a sI clathrate hydrate.15 In agreement with their analysis, singlecrystal X-ray structural analysis showed the Cl2 molecules are disordered between two symmetry distinct positions located close to the equatorial plane of the oblate large sI cages (see Figure 1a). At low temperatures, the disorder in molecule

a

The interaction with water transfers electron density from the water oxygen lone pair into the Cl−Cl σ* antibonding orbital and should lead to an elongation of the bond. However, in the clathrate hydrate phases, due to librational shortening, the Cl−Cl bond distances are shorter than the free Cl2 molecule bond length. See text for explanation. bExperimental value, ref 36. cBSSE corrected MP2/augcc-pVDZ calculations. dBSSE corrected MP2/6-311++G(d,p) calculations from BHHLYP/6-311++G(d,p) geometry optimized structure, ref 30.

Table 3. Bromine Bond Distances (Å) for the Free Molecule, the Water Complex with Water, and the Clathrate Hydrate Phases distance

d(H2O···X−X)

d(Br−Br) (free) d(Br−Br) (H2O···Br−Br) d(Br−Br) (sIII TA cage) d(Br−Br) (sIII TB cage) d(Br−Br) (sIII P cage) d(Br−Br) (sI mixed Cl2/Br2 T cage) d(H2O···Br−Br) (gas complex, calc.) d(H2O···Br−Br) (sIII TA cage) d(H2O···Br−Br) (sIII P cage) d(H2O···Br−Br) (sI Cl2/Br2 large cage)

2.28326a 2.338,b 2.319c 2.2737(12)−2.280(4) 2.268(3)−2.277(3) 2.270(3)−2.284(3) 2.235(4) 2.769,b 2.830c 2.98−3.174 3.15−3.23 2.75−2.85

a

Experimental value, ref 36. bBSSE corrected MP2/aug-cc-pVDZ calculations. cBSSE corrected MP2/6-311++G(d,p) calculations from BHHLYP/6-311++G(d,p) geometry optimized structure, ref 28.

Schofield and Jordan did not report the H2O···Br−Br distances in the cages. In all of their simulations, the small 512 cages of the clathrate hydrate phases were assumed empty. From the above discussions, chlorine and bromine clathrate hydrates present a number of unsolved questions. In the case of bromine hydrate, why is the structure unique to bromine with no other guests forming sIII? In the case of chlorine hydrate, what are the details of chlorine incorporation in the small cages? In this work we examine the detailed structures and guest−host interactions in chlorine and bromine hydrates as well as a mixed Cl2/Br2 hydrate using single crystal X-ray diffraction. We also summarize previous quantum chemical studies of halogen-water complexes to address some of these questions. The dihalogen−water interactions X−X···OH2 discussed in this work are described in terms of the well-known “halogen bonding” A−X···D that occurs between a Lewis acid halogen atom X (particularly Cl, Br, and I) of an A−X molecule and Lewis base atom D.29 Halogen bonding in the halogen−water

Figure 1. Chlorine guest molecule (a) disordered between the two symmetry-distinct positions large sI T cages and (b) in small sI D cages from single crystal X-ray diffraction. Some Cl···O distances are shown.

positions may be static; however, at higher temperatures disorder in the Cl2 positions in Figure 1 can be dynamic in nature. Since Pauling and Marsh considered the Cl2 molecule to be too large for the small cage, they did not analyze their diffraction data to study this feature of the hydrate. In our X-ray structure determination, the Cl2 guest molecules are not located in the center of the large sI T cages, and the smallest Cl···O distance found is 3.07 Å (see Figure 1a), which is much less than the sum of the O and Cl van der Waals radii (3.27 Å), indicating chlorine−water halogen bonding. 14178

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

Figure 2. Disordered positions of bromine guest molecules in (a) 14-sided TA-cages; (b) 14-sided TB-cages. Red lines represent short Br−Br···OH2 intermolecular distances less than 3.2 Å and yellow lines distances between 3.2−3.3 Å; (c) 15-sided P-cages.

evidence of bond elongation upon complexation with electron donating species when the chlorine molecule is not mobile. Although the short Cl−Cl···OH2 distance is evidence of halogen bonding between chlorine and water, due to the positions of the water molecules in lattice, the Cl−Cl···O angle for the closest distance are not 180°. To form the ideal 180° angle with chlorine, the lone pair electrons on a water molecule would need to point toward a chlorine atom. This could only occur if a water molecule rotates and breaks one of the water− water hydrogen bonds in the lattice. Furthermore, the Cl2 molecules are disordered in the cages, which precludes ideally directed halogen bonds for the guest in the cages. These observations show that halogen bonding can still stabilize water−halogen interactions, despite being multidirectional. Bromine Clathrate Hydrate. We previously studied the structure of bromine hydrate crystals of different stoichiometry with single-crystal X-ray structural analysis. At that time, extensive positional disorder of the bromine guests prevented us from localizing the positions of bromine molecules in the cages.23 Here we present results of structure refinement of the bromine positions in the cages. The bromine molecules are located in the large T and P cavities, while the small dodecahedral (D) cavities are vacant or partially occupied by O2 or N2 gases that are incorporated during crystallization in air. The bromine molecules are disordered over six independent crystallographic sites in the TA cavities (Figure 2a), eight sites in the TB cavities (Figure 2b), and six sites in the P cavities (Figure 2c). All bromine positions correspond to Wyckoff k sites of the P42/mnm space group with multiplicity of 16.

It is still a common practice in hydrate research to consider the van der Waals length of the guest molecule and the free diameter of the hydrate cage to see if the guest is a good fit for the cage under consideration.33 This “rule” would seem to rule out the presence of Cl2 guests in the small 512 sI cages. However, our experimental X-ray crystal structure analysis shows that there is a 32.5% probability of filling a small cage with a Cl2 guest, implying there are 0.65 Cl2 molecules in the small cages of each sI unit cell. In the structural determination of the small cage, the halogen−water oxygen interaction is even stronger than in the large cage as the shortest Cl···O distance is 2.90 Å (see Figure 1b). The bond length of the Cl2 guests in the sI clathrate hydrate (Table 2) from the X-ray structural analysis are determined to be 1.870 Å in the small D cages and 1.883−1.888 Å in the large T cages, all of which are shorter by 0.13−0.117 Å than the bond length of Cl2 in the gas phase (2.00 Å). In the Pauling and Marsh structure, the Cl−Cl bond is assigned a fixed length of 1.98 Å. In our case, the short Cl−Cl distances do not reflect short bond lengths in these encapsulated molecules but are related to the librational shortening and disorder of the Cl2 guest molecules in the sI cages.34 Quantum chemical calculations of Maity and co-workers,32 given in Table 2, show that the Cl−Cl bond length increases upon halogen bonding interaction of Cl2 with water, therefore the shortened Cl−Cl bond length in the single crystal X-ray structure of the clathrate hydrate is an artifact of the guest motion and disorder in the cages. The experimentally determined Cl−Cl bond length in dioxane-Cl2 cocrystals is 2.02 Å,35 which gives further 14179

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

Figure 3. (a) Chlorine (yellow) and bromine (brown) molecules in a binary sI chlorine and bromine hydrate. The bromine guests occupy the large cages (50.5%) and the chlorine guests occupy the large (49.5%) and small cages (59.3%). (b) Disordered water framework in the small cages of a binary sI chlorine and bromine hydrate. (c) Halogen bond between chlorine and water molecule in a distorted small cage of a binary sI chlorine and bromine hydrate. Elongated cage water−water edges are shown with blue lines. The minimum intermolecular Cl−Cl···OH2 distance in small cage is 2.80 Å.

for both bromine atoms in the smaller P cages. These distances are again considerably less than the Br···O van der Waals radius of 3.37 Å. The Br−Br···OH2 interaction in the bromine hydrate acts as a structure-directing synthon which gives rise to the unique sIII structure. The bond lengths of the Br2 guests in the sIII clathrate hydrate given in Table 2 are shorter by 0.015 to 0.005 Å than the bond length of Br2 in the gas phase. The heavy and larger Br2 guests are less mobile than Cl2 but still give rise to a small libration-induced shortening of bond length in the X-ray diffraction structure. Quantum chemical studies of Br− Br···OH2 complex28 show that the Br−Br bond length is elongated by electron donation from the water oxygen atom. Chlorine−Bromine Mixed Clathrate Hydrate. Pure bromine forms the unique tetragonal sIII clathrate hydrate structure, but a 1:1 mixture of bromine and chlorine forms cubic sI with bromine occupying large cages and chlorine molecules found in both small and large cages (see Figure 3a). For this mixture composition, the large (T) cages have roughly 50% probability of being occupied by Cl2 or Br2 (in Wyckoff l sites with multiplicity of 48) and small (D) cages are 59.3% occupied with Cl2 (in Wyckoff l sites with multiplicity of 48). The water framework in this sI clathrate hydrate is heavily disordered. Possible locations of water molecules in the D cage framework are shown in Figure 3b. Both large and small cages are distorted and short dihalogen-water interactions are found for both bromine and chlorine molecules. Figure 3c shows one of the possible disordered configurations of Cl2 in the distorted small cages where some cage edges are elongated from the normal 2.7−2.8 Å range to 3.0−3.1 Å. The smallest Br−Br···OH2 and Cl−Cl···OH2 distances in the cages are 2.75 and 2.70 Å, respectively. The Cl−Cl and Br−Br bond lengths in this mixed clathrate hydrate structure determined by X-ray diffraction show

In the TA cages, none of the bromine molecules have their molecular axis aligned toward the centers of the hexagonal faces, which could be due to the structure of the TA cavity, which is compressed along the axis perpendicular to the hexagonal faces and the tendency of bromine molecules to interact via halogen bonding with the water molecules at the face corners. Although the TA and TB 14-hedral cages have the same m symmetry, bromine molecules are arranged differently in them. Disordering of the bromine molecule in the P cavity is illustrated in Figure 2c, and the bromine molecules similarly point preferentially toward hexagonal face corners. The minimum intermolecular Br···O distance due to halogen bonding in a TB cage is 2.91 Å, significantly less than sum of van der Waals radii (3.37 Å). Short Br−Br···OH2 distances vary from 2.98 Å to 3.21 Å in the TA cavity, and from 3.15 Å to 3.26 Å in the P cavity. As the case for the Cl2 clathrate hydrate, we believe the large degree of disorder is related to the dynamical motions of the Br2 molecules in the cages. Molecular dynamics simulations of the sIII bromine hydrate phase would clarify the situation. The librational shortening of the Br2 molecule in the cages (discussed below) is another evidence for the motions of these guests in the cages. Using standard dimensions, the van der Waals length of the Br2 molecule is 5.98 Å (2.28 + 2 × 1.85) and based on van der Waals radius alone Br2 would be considered too large for both the 512 D and 51262 T (5.8 Å diameter) cages. Yet in the sIII hydrate, Br2 occupies T and P cages, which is due to the halogen bonding and shorter Br−Br···OH2 distances than predicted by van der Waals radii. The sII large 16-sided cage (H = 51264) with a diameter of 6.4 Å is too large to give close contacts between different waters oxygen atoms and both bromine atoms in Br2. This could suggest why the sIII hydrate structure is favored for bromine, since close Br−Br···OH2 contacts of 3.15 −3.26 Å, are observed 14180

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

occur if (as observed in the single-crystal structure) water molecules are displaced from their ideal positions in the clathrate hydrate. The variety of interactions between water and guest molecules in the clathrate hydrate can include hydrogen bonding, where water can act as a proton donor or acceptor (Brønsted acid and or base), and halogen bonding, where water acts as an electron pair donor (Lewis base). The water−halogen interactions do not preclude the formation of ordered clathrate hydrate phases, but can affect the nature of the hydrate formed, for example, sIII for Br2 or a disordered sI for the Br2/Cl2 mixture). Experimental Single-Crystal Structure Details. Single crystal X-ray diffraction data were measured on a Bruker Apex 2 Kappa diffractometer at 100 K, using graphite monochromatized Mo Kα radiation (λ = 0.71073 Å). The unit cell was determined from randomly selected reflections obtained using the Bruker Apex2 automatic search, center, index, and leastsquares routines. Integration was carried out using the program SAINT, and an absorption correction was performed using SADABS.37 The crystal structures were solved by direct methods and the structure was refined by full-matrix leastsquares routines using the SHELXTL program suite.38 All atoms were refined anisotropically. Hydrogen atoms on guest molecules were placed in calculated positions and allowed to ride on the parent atoms. The crystallographic information files for all three structures are given in the Supporting Information. Chlorine Clathrate Hydrate. Chemical formula: Cl13.3O46H92.[Cl2·6.92H2O], M = 209.02, Z = 1, crystal size 0.5 × 0.4 × 0.3 mm3, cubic space group Pm3̅n, a = 11.9710(1)Å, V = 1715.50(2) Å3, T = 173(2) K, ρ = 1.619 g·cm−3, 17919 reflections measured, 429 unique 360 [I > 2σ(I)], final R indices [I > 2σ(I)] R1 = 0.0152, wR2 = 0.0366, R indices (all data) R1 = 0.0209, largest differential peak and hole 0.08 and −0.04 e·Å−3, respectively. Bromine Clathrate Hydrate. Chemical formula: Br38.27O172H344.[Br2·8.99H2O], M = 306.13, Z = 1, crystal size 0.4 × 0.35 × 0.25 mm3, tetragonal space group P4(2)/mnm, a = 23.0436(9) Å, c = 12.0745(7) Å, V = 6411.7(5) Å3, T = 173.0(1) K, ρ = 1.595 g·cm−3, 24788 reflections measured, 2549 unique, 1868 [I > 2σ(I)], final R indices [I > 2σ(I)] R1 = 0.0384, wR2 = 0.0831, R indices (all data) R1 = 0.0740, largest differential peak and hole 0.66 and −0.27 e·Å−3, respectively. Chlorine and Bromine Mixed Clathrate Hydrate. Chemical formula: Cl8.31Br6.06O46H92, M = 352.42, Z = 1, crystal size 0.3 × 0.2 × 0.1 mm3, cubic space group Pm3̅n, a = 11.9619(4) Å, V = 1711.6(1) Å3, T = 173.0(1) K, ρ = 1.710 g·cm−3, 19633 reflections measured, 500 unique 384 [I > 2σ(I)], final R indices [I > 2σ(I)] R1 = 0.0261, wR2 = 0.0829, R indices (all data) R1 = 0.0431, largest differential peak and hole 0.19 and −0.14 e·Å−3, respectively.

librational shortening of 0.139−0.135 and 0.048 Å, respectively. This gives indication that these guests are mobile in the large and small cages and the crystallographic disorder is of dynamic nature.

■

CONCLUSIONS The Cl2, Br2, and their mixed clathrate hydrates show guestwater interactions consistent with halogen bonding. From considerations of van der Waals radii we would not expect Cl2 to occupy sI clathrate hydrate small D cages, but in the pure Cl2 hydrate 59% experimental occupancy of the small cages is observed. Measured Cl···O distances in the all sI clathrate hydrate cages are smaller than predicted by the sum of van der Waals radii. This is not unexpected as gas phase water and Cl2 complexes show halogen bonding with a binding energy of ∼2.8 kcal/mol30 and their formation is accompanied by the elongation of the Cl−Cl bond through the transfer of electron density from the oxygen atom lone pair into the Cl2 σ* antibonding orbital. A similar picture holds for the Br2 hydrate where Br···O distances in the all sIII clathrate hydrate cages are smaller than predicted by the sum of van der Waals radii. Gas phase Br2−water complexes have a larger binding energy of 3.9 kcal/mol,30 and their formation also is accompanied by Br−Br elongation. Two features of the clathrate hydrate phase complicate this picture. In a perfect defect-free clathrate hydrate cage, all water molecules have four hydrogen bonds and there are no lone pair electrons available for interactions with Br2 and Cl2. However, due to the significant H2O···Cl2 and H2O···Br2 interaction energies, it is plausible that water molecules in the clathrate hydrate cages rotate to allow lone-pair electrons to interact with the Cl2 and Br2 guest, forming a Bjerrum defect in the hydrogen bonded water−water hydrate lattice. These interactions would make the H2O···X distances shorter than expected by simple considerations of the van der Waals radii. The fact that some Cl2 molecules are observed in the small cages, where according to considerations of the van der Waals radii they should not fit, indicates halogen bonding interactions between Cl2 and the water lattice, perhaps accompanied by water rotations and formation of Bjerrum defects and lattice distortions. Differences between the clathrate hydrate phases of the Cl2 and Br2 could be related to the larger binding energy or the size of Br2 compared to Cl2. The bond length of the halogen guests are given in Tables 2 and 3. Librational shortening gives apparent Cl−Cl bond lengths from the single crystal X-ray measurements in the clathrate hydrate phases which are consistently shorter than the free Cl−Cl bond in the gas phase and the calculated Cl2···(H2O)n clusters. The Br−Br bond lengths are also shortened compared to the gas phase Br−Br bond length. The libration shortening gives evidence of the dynamic motions of the guests in the cages. In water-halogen clusters, as the number of water molecules increases, the X−X bond length increases and X···O distance decreases. The X···O distances in the water clusters are generally shorter than the contact distances of the halogens with water in the clathrate hydrate cages as linear X−X···OH2 chemical interactions in the clathrate are not possible due to lattice positions of the waters. In the mixed Cl2/Br2 sI clathrate hydrate, the minimum Cl···O and Br···O distances in the large cages are short, 2.80 and 2.876 Å, respectively. These short distances are consistent with strong halogen-water chemical interactions which can only

■

ASSOCIATED CONTENT

* Supporting Information S

Crystallographic information files for single-crystal X-ray structures of the Cl2, Br2, and mixed Cl2/Br2 clathrate hydrates. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. 14181

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182

The Journal of Physical Chemistry C

Article

Author Contributions

(27) Legon, A. C.; Thumwood, J. M. A.; Waclawik, E. R. Chem. Eur. J. 2002, 8, 940. (28) (a) Bernal-Uruchurtu, M. I.; Hernández-Lamoneda, R.; Janda, K. C. J. Phys. Chem. A 2009, 113, 5496−5505. (b) Franklin-Mergarejo, R.; Rubayo-Soneira, J.; Halberstadt, N.; Ayed, T.; Bernal-Urchurtu, M. I.; Hernández-Lamoneda, R.; Janda, K. C. J. Phys. Chem. A 2011, 115, 5983. (29) (a) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Acc. Chem. Res. 2005, 38, 386−395. (b) Politzer, P.; Lane, P.; Concha, M. C.; Ma, Y.; Murray, J. S. J. Mol. Model. 2007, 13, 305−311. (30) (a) Zhou, P.; Lv, J.; Zou, J.; Tian, F.; Shang, Z. J. Struct. Biol. 2010, 169, 172−182. (b) Ibrahim, M. M. J. Comput. Chem. 2011, 32, 2564−2574. (c) Kolár,̌ M.; Hobza, P. J. Chem. Theory Comput. 2012, 8, 1325−1333. (d) Ř ezác,̌ J.; Riley, K. E.; Hobza, P. J. Chem. Theory Comput. 2012, 8, 4285−4292. (e) Jorgensen, W. L.; Schyman, P. J. Chem. Theoy. Comput. 2012, 8, 3895−3901. (31) The use of halogen bonding to address interactions in which halogen atoms function as electron-donors is conceptually misleading and contrasts with the clear tendency, well-documented in the literature, to name such −X···H−Y interactions differently (i.e., hydrogen bonds). See http://www.iupac.org/web/ins/2009-032-1100. (32) (a) Pathak, A. K.; Mukherjee, T.; Maity, D. K. J. Chem. Phys. 2006, 124, 024322. (b) Pathak, A. K.; Mukherjee, T.; Maity, D. K. J. Phys. Chem. A 2008, 112, 744−751. (33) Sloan Jr., E. D.; Koh, C. A. Clathrate Hydrates of Natural Gases, 3rd ed.; CRC Press: Boca Raton, FL, 2007. (34) (a) Cruikshank, D. W. J. Acta Crystallogr. 1956, 9, 757. (b) Jones, G. P. Chem. Soc. Rev. 1984, 13, 157. (35) Hassel, O.; Strømme, K. O. Acta Chem. Scand. 1959, 13, 1775. (36) Barrow, R. F.; Clark, T. C.; Coxon, J. A.; Yee, K. K. J. Mol. Spectrosc. 1974, 51, 428−449. (37) Sheldrick, G. M. SADABS, version 2.03; University of Gottingen: Germany, 2002. (38) Sheldrick, G. M. SHELXTL, version 6.10; Bruker AXS, Inc.: Madison, WI, 2000.

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS The authors thank the National Research Council of Canada for support.

■

REFERENCES

(1) Jeffrey, G. A. In Comprehensive Supramolecular Chemistry; Atwood, J. L., Davies, J. E. D., MacNicol, D. D., Vögtle, F., Eds.; Pergamon: Oxford, U.K., 1996; Vol. 6, p 757. (2) Alavi, S.; Udachin, K. A.; Ratcliffe, C. I.; Ripmeester, J. A. Clathrate Hydrates. In Supramolecular Chemistry: From Molecules to Nanomaterials; John Wiley & Sons, Ltd.: Chichester, U.K., 2011. (3) Ripmeester, J. A. Ann. N.Y. Acad. Sci. 2000, 912, 1−16. (4) van der Waals, J. H.; Platteeuw, J. C. Adv. Chem. Phys. 1959, 2, 1. (5) Davidson, D. W. In Water, A Comprehensive Treatise; Franks, F., Ed.; Plenum Press: New York, 1973; Vol. 2, pp 115−234. (6) (a) Mootz, D.; Oellers, E.-J.; Wiebcke, M. J. Am. Chem. Soc. 1987, 109, 1200. (b) Desmedt, A.; Stallmach, F.; Lechner, R. E.; Cavagnat, D.; Lassègues, J.-C.; Guillaume, F.; Grondin, J.; Gonzales, M. A. J. Chem. Phys. 2004, 121, 11916. (7) (a) Davidson, D. W.; Ripmeester, J. A. In Inclusion Compounds; Atwood, J. L., Davies, J. E. D., MacNicol, D. D., Eds.; Academic Press: London, 1984; Vol. 3. (b) Gough, S. R.; Hawkins, R. E.; Morris, B.; Davidson, D. W. J. Phys. Chem. 1973, 77, 2969. (8) Williams, K. D.; Devlin, J. P. J. Mol. Struct. 1997, 416, 277. (9) (a) Koga, K.; Tanaka, H.; Nakanishi, K. Mol. Simul. 1994, 12, 241. (b) Koga, K.; Tanaka, H.; Nakanishi, K. Mol. Simul. 1996, 16, 151. Koga, K.; Tanaka, H. J. Chem. Phys. 1996, 104, 263. (10) (a) Alavi, S.; Susilo, R.; Ripmeester, J. A. J. Chem. Phys. 2009, 130, 174501. (b) Alavi, S.; Udachin, K. A.; Ripmeester, J. A. Chem. Eur. J. 2010, 16, 1017. (11) (a) Buch, V.; Devlin, J. P.; Monreal, I. A.; Jagoda-Cwiklik, B.; Aytemiz-Uras, N.; Cwiklik, L. Phys. Chem. Chem. Phys. 2009, 11, 10245. (b) Monreal, I. A.; Cwiklik, L.; Jagoda-Cwiklik, B.; Devlin, J. P. J. Phys. Chem. Lett. 2010, 1, 290. (12) Shin, K.; Alavi, S.; Udachin, K. A.; Ripmeester, J. A. Proc. Natl. Acad. Sci. U.S.A. 2012, 109, 14785−14790. (13) Davy, H. Philos. Trans. R. Soc. 1811, 101, 155. (14) Faraday, M. Quart. J. Sci. 1823, 15, 71. (15) Pauling, L.; Marsh, R. Proc. Natl. Acad. Sci. U.S.A. 1952, 38, 112−119. (16) Glew, D. N.; Rath, N. S. J. Chem. Phys. 1966, 44, 1710. (17) Anwar-Ullah, S. J. Chem. Soc. 1932, 1172. (18) Cady, G. J. Phys. Chem. 1983, 87, 4441; J. Phys. Chem. 1985, 89, 3302. (19) (a) Löwig, C. Mag. Pharm. 1828, 23, 12. (b) Löwig, C. Ann. Chim. Phys., Ser. 2 1829, 42, 113. (20) von Stackelberg, M.; Müller, H. R. Z. Elektrochem. 1954, 58, 25. (21) Allen, K. W.; Jeffrey, G. A. J. Chem. Phys. 1963, 38, 2304. (22) Dyadin, Y. A.; Aladko, L. S. Zh. Struct. Khim. 1977, 18, 51. (23) Udachin, K. A.; Enright, G. D.; Ratcliffe, C. I.; Ripmeester, J. A. J. Am. Chem. Soc. 1997, 119, 11481−11486. (24) (a) Kerenskaya, G.; Goldschleger, I. U.; Apkarian, V. A.; Janda, K. C. J. Phys. Chem. A 2006, 110, 13792−13798. (b) Goldschleger, I. U.; Kerenskaya, G.; Janda, K. C.; Apkarian, V. A. J. Phys. Chem. Lett. 2008, 112, 787−798. (25) Schofield, D. P.; K. D. Jordan, K. D. J. Phys. Chem. A 2009, 113, 7431−7438. (26) Hassel, O. Nobel Lectures, Chemistry 1963−1970; Elsevier: Amsterdam, 1972. 14182

dx.doi.org/10.1021/jp402399r | J. Phys. Chem. C 2013, 117, 14176−14182