Hydrogen-Bonded Helices for Anion Binding and Separation Radu Custelcean,*,† De-en Jiang,† Benjamin P. Hay,† Wensui Luo,‡ and Baohua Gu‡ Chemical Sciences DiVision and EnVironmental Sciences DiVision, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831
CRYSTAL GROWTH & DESIGN 2008 VOL. 8, NO. 6 1909–1915
ReceiVed February 4, 2008; ReVised Manuscript ReceiVed March 14, 2008
ABSTRACT: Herein we report the competitive crystallization of urea-functionalized hydrogen-bonded helical frameworks as a new approach to separating anions from aqueous mixtures. N,N′-Bis(m-pyridylurea) (1) containing orthogonal pyridine and urea hydrogen-bonding functionalities forms upon monoprotonation with 1 equiv of HX acids (X ) Cl-, Br-, I-, NO3-, and ClO4-) an isomorphous series of crystalline hydrogen-bonded helices assembled by pyridinium · · · pyridine hydrogen bonds, with the urea functional groups binding the anions through chelate hydrogen bonding. The helices are further connected in the crystals by CH · · · Xand pyridinium · · · X- interactions, as well as π-stacking interactions. Competitive crystallization experiments and lattice energy calculations of the 1 · HX crystals showed the solvation-based Hofmeister bias that typically dominates anion separation selectivities from water was attenuated, but not completely overturned. The observed selectivity is apparently a result of the relatively soft and unspecific hydrogen-bonding environment around the anions in the crystals, combined with the high flexibility of the helices, which expand or contract as necessary to accommodate each anion. Introduction Hydrogen bonds play critical roles in nature, from governing the structure and properties of water to ensuring the cohesion and functionality of complex biological structures.1 It is therefore of little surprise that researchers have long attempted to utilize these interactions for the deliberate assembly of materials and nanostructures with desired architectures and functions.2 An important step in this process is the judicious selection of a persistent H-bonding synthon that will act as “glue” holding the molecular components together.3 The design is considered successful if all synthons are utilized in the assembly of the framework according to their original designation. If a particular function is sought, the starting building units also need to be functionalized with appropriate functional groups that will manifest the desired property in the final material. A potential conflict arises when these functional groups have themselves H-bonding abilities, as they may interfere with the framework assembly by disrupting the “expected” H-bonded motifs, or may be neutralized through self-association. For these reasons, functionalization of crystalline frameworks with H-bonding groups has rarely been achieved, and the few successful attempts to date are limited to metal-organic frameworks or coordination polymers where the H-bonding and metal coordination functionalities can be more readily separated.4 Along this line, we have recently designed and synthesized metal-organic frameworks (MOFs) functionalized with H-bonding groups for anion recognition and separation.4a,5 We were particularly successful in functionalizing MOFs with anion-chelating urea groups as exemplified by the prototype N,N′-bis(m-pyridylurea) (1),5d,e,6 in which the pyridine and urea functional groups were found to operate independently as framework builder and anion binder, respectively. Here we extend this approach to H-bonded frameworks built from the same linker 1, in which both the framework assembly and anion binding roles are played by hydrogen bonds. * To whom correspondence should be addressed. E-mail: custelceanr@ ornl.gov. † Chemical Sciences Division. ‡ Environmental Sciences Division.
Figure 1. The concept of orthogonal H-bond functionalities applied to the assembly of hydrogen-bonded frameworks functionalized with urea groups for anion binding.
The key to the design of hydrogen-bonded frameworks functionalized with anion-binding groups lies in employing orthogonal H-bond functionalities as illustrated in Figure 1. Our approach is based on Etter’s empirical H-bond rules,7 which postulate that (1) “all good proton donors and acceptors are used in hydrogen bonding”; (2) “six-membered-ring intramolecular hydrogen bonds form in preference to intermolecular hydrogen bonds”; and (3) “the best proton donors and acceptors remaining after intramolecular hydrogen-bond formation form intermolecular hydrogen bonds to one another”. It thus follows that upon monoprotonation, 1 is set up to form H-bonded chains by utilizing the best proton donor and acceptor, which are the pyridinium and the pyridine groups, respectively. The urea carbonyl group is expected to form six-membered-
10.1021/cg800137e CCC: $40.75 2008 American Chemical Society Published on Web 06/04/2008
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Figure 3. Interdigitation of helices in 1 · HX resulting in anion intercalation between pyridinium groups. X ) Cl- (a), Br- (b), I- (c), NO3- (d), and ClO4- (e). Figure 2. Binding of anions (ball-and-stick model) by H-bonded helices (stick model) functionalized with urea functional groups in crystals of 1 · HX, as determined by X-ray crystallography. X ) Cl- (a), Br- (b), I- (c), NO3- (d), and ClO4- (e).
ring intramolecular H-bonds involving the ortho C-H donors,5e,8 which are activated by the electron-withdrawing pyridines, similarly to the prototype bis(p-nitrophenyl)urea originally studied by Etter.9 The CdO group is thus less available to engage in intermolecular H-bonds with the NH donors, as typically observed in N,N′-disubstituted ureas that have a strong propensity to form H-bonded tapes.2a As a result, the urea NH donors should remain available to bind the anions. In the following pages we demonstrate that, conforming to these design principles, 1 forms in the presence of 1 equiv of HX acids (X ) Cl-, Br-, I-, NO3-, and ClO4-) an isomorphous series of hydrogen-bonded helices10 assembled by pyridinium · · · pyridine H-bonds, with the urea functional groups binding the anions through chelate H-bonding. The occurrence of a common structural motif throughout this series offered a rare opportunity to evaluate the factors determining anion separation selectivity in hydrogen-bonded frameworks, which was investigated here by competitive crystallization experiments5,11 from aqueous anionic mixtures, and theoretical calculations. Results and Discussion Crystallization of 1 in the presence of 1 equiv of HX (X ) Cl-, Br-, I-, NO3-, and ClO4-) yielded an isomorphous series of 1 · HX salts, as revealed by single-crystal X-ray diffraction. While the urea CdO acceptors were locked into six-memberedring intramolecular H-bonds with the pyridine ortho C-H donors, monoprotonation of 1 to one of its two pyridine groups induced the self-assembly into 21-symmetrical helices in the solid state, held together by pyridinium · · · pyridine H-bonds (Figure 2). Once these strongest H-bond donors and acceptors were neutralized, the urea NH donor groups remained available to bind the anions through chelate hydrogen bonding. Thus, as expected from Etter’s H-bond rules, all good proton donors and acceptors were hierarchically involved in hydrogen bonding,2i,j thereby ensuring the separation of the H-bond functionalities into two noninterfering orthogonal tasks: framework assembly and anion binding. The helices are further connected into layers
in the (1 0 0) plane by interdigitation, with the anions from one helix intercalated between the pyridinium groups of the neighboring helices (Figure 3). The helical pitch, coinciding with the unit cell b-axis, adjusts appropriately to accommodate the sizes and shapes of the different anions, through twisting about the pyridinium · · · pyridine H-bonds that act as flexible “hinges”. Alternating stacking of the pyridine and pyridinium groups also contributes to the cohesion between helices. While there is little variation in the helical pitch among the halide structures (from 7.70 to 7.79 Å), the helices shrink considerably to b ) 7.28 Å to accommodate the flatter NO3- and expand to b ) 7.84 Å to fit the bulkier ClO4-. Finally, the helices assemble along the crystallographic a direction via C-H · · · X hydrogen bonding involving the pyridinium and pyridine CH donors,12 to complete the crystal packing (Figure 4). In addition to the adjustable helical axis b, the flexibility of these crystals is also evident in the variation of the a and c unit cell parameters, which adjust accordingly to fit the size of the anions, increasing monotonically from Cl- to ClO4- (Table 4). Overall, the unit cell volume increases by 16.8% when going from chloride to perchlorate. This flexibility is presumably critical in maintaining the same crystal structure throughout the series, despite the wide variations in the anions’ sizes and shapes. Furthermore, despite the significant difference in anions’ basicity (most extremely manifested between Cl- and ClO4-), all anions display very similar coordination environments. Thus, as illustrated in Figure 5, each anion is H-bonded by a chelating urea group (I), two pyridinium groups12 (ortho CH (II), and bifurcated metha-para CH (III)) and a pyridine group (metha or bifurcated ortho-meta CH (IV)). Additionally, two pyridinium groups flank the anions with their π rings. These interactions (V-VI) are however different from the so-called anion-π interactions, as the rather long contacts with the aromatic rings involve C atoms from the periphery rather than the ring’s centroids.13 The geometrical parameters of the interactions I-VI in the five crystal structures are tabulated in Table 1. To assess the relative energetic contributions of the interactions I-VI to anion binding in these crystals, we performed electronic structure calculations using B3LYP/6-311+G** on the prototype system 1 · HCl. Thus, a fragment representing each
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Figure 4. Packing of helices in 1 · HX viewed along the helical axis b. X ) Cl- (a), Br- (b), I- (c), NO3- (d), and ClO4- (e). Table 1. Observed Anion Coordination Parameters (Å, °) in 1 · HX (X ) Cl-, Br-, I-, NO3-, ClO4-) and Calculated Chloride Binding Energies (kJ/mol) in 1 · HCl (shown in bold)a for Interactions I-VI 1 · HCl I
2.28, 2.40, II 2.77, III 2.81, 3.17, IV 2.91, V 3.75 VI 3.49
Figure 5. Anion coordination environments in 1 · HX. X ) Cl- (a), Br- (b), I- (c), NO3- (d), and ClO4- (e). H-bonds are shown as black dashed lines and X- · · · C(pyridine) as green dashed lines.
binding motif was isolated from the crystal using the coordinates from the X-ray structure, and then the fragment was optimized under constraints to maintain the same configuration as found in the crystal. A chloride anion was subsequently positioned relative to the fragment to have the spatial relationship as close as possible to that observed in the crystal. A single point energy of the complex was finally computed at the B3LYP/6-311+G** level,14 and the electronic binding energy was calculated as ∆E ) E(complex) - E(fragment) - E(Cl-) and listed in Table 1. As expected, such interactions between charged species in the gas phase are much stronger than typically expected in a crystal, where all these interactions (and other long-range ones) occur simultaneously and in an overall equilibrium. Nevertheless, this methodology allows an estimation of the relative importance of the individual interactions to the overall anion stabilization in the crystal.
163.5 157.6 162.8 135.1 118.8 142.7
1 · Br -422 2.46, 2.53, -279 2.87, -292 2.93, 3.33, -235 3.02,
1 · HI
164.5 160.4 162.8 137.0 119.1 140.5
-281 3.70 -281 3.51
2.69, 2.73, 3.06, 3.13, 3.45, 3.21, 3.76 3.60
166.9 163.1 162.2 136.5 121.8 138.0
1 · HNO3 1.97, 1.98, 2.35, 2.48, 2.77, 2.34,
173.2 167.2 133.1 130.3 117.4 157.6
3.47 3.25
1 · HClO4 2.06, 2.06, 2.37, 2.65, 2.84, 2.58, 2.77, 3.11, 3.08,
165.0 164.2 161.3 124.4 117.2 125.2 116.5 3.20 3.24
a Fragments were chosen such that each of the binding motifs contained one positive charge. For the urea interaction, I, the fragment was protonated N,N′-bis(m-pyridylurea). For the interactions with pyridine and pyridinium moieties, II-VI, a urea substituent terminated with an NH2 group was attached to the binding arene and either a pyridine or pyridinium cation was added forming a N-H---N hydrogen bond with the binding arene. Cartesian coordinates for each of the fragments are provided as Supporting Information.
Table 2. Energetics of Anion Binding and Separation in 1 · HX. All Energies are in kJ/mol 1 · HCl
1 · Br
1 · HNO3
1 · HI
1 · HClO4
∆Elatt -1230.1 -1205.5 -1196.3 -1178.8 -1151.8 ∆G0h (X-)b -347 -321 -306 -275 -214 c ∆Eexch (g) -78.3 -53.7 -44.5 -27.0 0 ∆G0exch (aq) (comp)d 54.7 53.3 47.5 34.0 0 ∆G0exch (aq) (exp)e 7.8 4.0 3.3 0 1.5 a
a Lattice energies: ∆Elatt ) E(1 · HX)(cryst) - E(1 · H+)(g) - E(X-)(g). Experimental anions’ Gibbs energies of hydration. c Calculated energies for perchlorate exchange reactions in the gas phase: ∆Eexch (g) ) ∆Elatt(1 · HX) - ∆E latt(1 · HClO4). d Calculated free energies of reaction for perchlorate exchange in water: ∆G0exch (aq) ) ∆Eexch (g) ∆G0h(X-) + ∆G0h(ClO4-). e Estimated experimental free energies of reaction for iodide exchange in water. b
Our calculations found that although the urea fragment is the strongest anion binder, the other fragments, with binding energies ranging between 55% and 70% of the corresponding
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Figure 6. Energetics of anion binding and exchange in 1 · HX. Blue ) experimental Gibbs free energy of dehydration (-∆G0h) relative to ClO4-; red ) calculated energies for perchlorate exchange reactions in the gas phase: ∆Eexch (g) ) ∆Elatt(1 · HX) - ∆Elatt(1 · HClO4); green ) calculated free energies of reaction for perchlorate exchange in water: ∆G0exch (aq) ) ∆Eexch (g) - ∆G0h(X-) + ∆G0h(ClO4-); black ) estimated experimental free energies of reaction for iodide exchange in water. Table 3. Iodide Separation Factors r(I-/X) ) [(mol I-)/(mol X)]crystal/[(mol I-)/(mol X)]solution for the Pairwise Competition Experiments anion (X) -
Cl BrNO3ClO4-
crystallization yield (%)
[(I-)/(X)]crys
[(I-)/(X)]sol
R(I-/X)
45 12 15 39
13.3 4.6 3.5 1.6
0.6 0.92 0.91 0.90
22.2 5 3.8 1.8
binding by the urea, make significant contributions toward chloride binding, and in fact, together they provide most of the anion’s stabilization within the binding cavity. This is in line with our prior calculations, showing that interactions between C-H groups or π systems of electron-deficient arenes and anions can be relatively strong.13,15 The structural persistence of the hydrogen-bonded helices in 1 · HX offers a rare opportunity to interrogate the factors governing anion selectivity in crystallization of hydrogenbonded frameworks. Since all five structures are virtually identical, the observed selectivity can be directly correlated with the anions’ coordination environment in the crystals. Such structure-selectivity relationships are not easily interpreted in the vast majority of anion-binding crystalline systems, as the great variance in crystal packing from one anion to another often causes an unpredictable change in solubility and thereby selectivity.5b To assess the anion selectivity in 1 · HX, we set up a competition experiment in which 1 equiv of 1 was crystallized from an aqueous mixture containing 1 equiv of each Cl-, Br-, I-, NO3-, and ClO4-, added as sodium salts, in the presence of H2SO4 as a source of protons (H2SO4 alone was found not to form any crystals with 1 under these conditions). The optimum amount of H+ was found to be about 1.5 equiv, with lower amounts resulting in decreased crystallization yields, and higher than 2 equiv resulting in no crystallization. The crystallized solid was analyzed by single-crystal X-ray diffraction and ion chromatography to evaluate its anionic composition. The X-ray analysis found the same structure as in the individual 1 · HX salts, but with intermediate unit cell parameters (comp cryst in Table 5), suggesting the formation of a solid solution including different amounts of anions. While only the most abundant I- and ClO4- were included in the crystal structure refinement, the ionic analysis of the bulk solid indicated in
addition to these two anions the presence of NO3- and Br-, with a molar ratio of I-/ClO4-/NO3-/Br- ) 14/7/3.4/1. Pairwise competitive crystallizations of I- with each of the other anions were also performed to independently determine the anion inclusion selectivity relative to iodide by single-crystal X-ray diffraction. As shown in Table 5, the binary 1 · H(I)y(X)1-y (X ) Cl-, Br-, NO3-, ClO4-) crystals thus obtained are all solid solutions containing Ias the dominant anion and variable proportions of the other anions, from which the following selectivity could be inferred: I-/ClO4-/ NO3-/Br-/Cl- ) 13.3/8.3/3.8/2.9/1. Thus, both methods established the same relative selectivity of I- > ClO4- > NO3- > Br> Cl-, which largely follows the Hofmeister series,5b except the order of iodide and perchlorate is reversed. To gain insight into the observed anion selectivity, we performed lattice energy calculations on the 1 · HX crystals using density functional theory at the Perdew-Berke-Ernzerhof (PBE) level with periodic boundary conditions and planewave bases. The calculated lattice formation energies from the gas phase ions were calculated as ∆Elatt ) E(1 · HX)(cryst) - E(1 · H+)(g) - E(X-)(g), and are listed in Table 2, along with anions’ experimental Gibbs energies of hydration (∆G0h(X-)).16 The free energy ∆G0exch (aq) for the hypothetical perchlorate exchange in water (reaction 4) was calculated by combining reactions 1–3. Thus, the overall process was conveniently broken down into three elementary steps: gas phase perchlorate exchange (1), dehydration of X(2), and hydration of ClO4- (3). It thus follows that ∆G0exch 0 0 0 (aq) ) ∆G exch (g) - ∆G h(X ) + ∆G h(ClO4 ), where the free energy for perchlorate exchange reactions in the gas phase was approximated as ∆G0exch (g) ) ∆Eexch (g) ) ∆Elatt(1 · HX) - ∆Elatt(1 · HClO4). 1·HClO4(cryst) +X(g) f 1·HX(cryst) + ClO4(g)
(1)
X(aq) f X(g)
(2)
ClO4(g) f ClO4(aq)
(3)
1·HClO4 (cryst) + X(aq) f 1·HX(cryst) + ClO4(aq)
(4)
Figure 6 depicts graphically the predicted anion selectivities (green triangles), plotted as ∆G0exch (aq) vs the inverse anion radii (1/r). This representation is a convenient way to visualize anion selectivity results, as it allows comparison against the Hofmeister series, represented here as the monotonic increase of the -∆G0h (blue diamonds) with decreasing anion radii. The calculated ∆G0exch (g), on the other hand, represented as red squares, shows a monotonic decrease from ClO4- to Cl-, which follows the decrease in the anion size and the corresponding increase in charge density. As illustrated in Figure 6, our calculations predict the attenuation of the Hofmeister bias, evidenced by flattening of the ∆G0exch (aq) (calcd) curve (green line) relative to the ∆G0h one (blue line). This is a consequence of the fact that crystallization selectivity from water is the sum of two monotonic but opposite biases: the anions’ free energy of dehydration (blue line), which increases with 1/r, and the lattice energy (red line), which decreases (becomes more negative) with 1/r. For comparison, the experimental crystallization selectivities (∆G0exch (aq) (exp)), estimated from the pairwise competition experiments (Table 3), are shown as black circles. The good qualitative agreement between the calculated and experimental selectivities is evident, as both series show attenuation of the Hofmeister bias. However, compared to theory, the experimental results show a more pronounced attenuation of the solvation bias, and even a small peak selectivity for I-. This difference
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Crystal Growth & Design, Vol. 8, No. 6, 2008 1913
Table 4. Crystallographic Data for 1 · HX (X ) Cl-, Br-, I-, NO3-, ClO4-) formula M crystal size [mm] crystal system space group a [Å] b [Å] c [Å] R [°] β [°] γ [°] V [Å3] Z T (K) Fcalcd [g cm-3] 2θmax [°] reflns collected independent reflns parameters R1, wR2 (I > 2σ(I)) GOF
1 · HCl
1 · Br
1 · HI
1 · HNO3
1 · HClO4
C11H11ClN4O 250.69 0.5 × 0.3 × 0.11 monoclinic P21/c 11.6266(10) 7.7488(7) 13.1143(11) 90 102.664(2) 90 1152.75(17) 4 173 1.444 56.58 8000 2777 154 0.0355, 0.0963 1.046
C11H11BrN4O 295.15 0.34 × 0.24 × 0.04 monoclinic P21/c 11.8359(14) 7.6971(9) 13.3409(16) 90 102.711(2) 90 1185.6(2) 4 173 1.654 56.70 14585 2951 154 0.0293, 0.0758 1.034
C11H11IN4O 342.14 0.39 × 0.16 × 0.04 monoclinic P21/c 12.1613(6) 7.7888(4) 13.5764(6) 90 102.6820(10) 90 1254.61(11) 4 173 1.811 56.58 10498 3114 174 0.0200, 0.0496 1.033
C11H11N5O4 277.25 0.27 × 0.26 × 0.08 monoclinic P21/c 12.5138(6) 7.2841(3) 13.7990(6) 90 103.2970(10) 90 1224.08(9) 4 173 1.504 56.62 12564 3044 181 0.0462, 0.1097 1.060
C11H11ClN4O5 314.69 0.34 × 0.27 × 0.08 monoclinic P21/c 12.5996(10) 7.8381(6) 13.9893(11) 90 102.969(2) 90 1346.30(18) 4 173 1.553 56.58 16599 3351 190 0.0643, 0.1885 1.058
Table 5. Crystallographic Data of the Binary Crystals 1 · H(I)y(X)1-y (X ) Cl-, Br-, NO3-, ClO4-) Obtained from 1:1 I-/X- Mixtures, and of the Mixed Crystals from the Competitive Crystallization Experiment (comp cryst)
formula M crystal size [mm] crystal system sp group a [Å] b [Å] c [Å] R [°] β [°] γ [°] V [Å3] Z T (K) Fcalcd [g cm-3] 2θmax [°] reflns collected independent reflns parameters R1, wR2 (I > 2σ(I)) GOF
1 · H(I)0.93(Cl)0.07
1 · H(I)0.82(Br)0.18
1 · H(I)0.78(NO3)0.22
1 · H(I)0.61(ClO4)0.39
comp cryst
C11H11Cl0.07I0.93N4O 335.74 0.35 × 0.34 × 0.05 monoclinic P21/c 12.1661(10) 7.7894(7) 13.5686(11) 90 102.692(2) 90 1254.43(18) 4 173 1.778 56.58 15517 3119 155 0.0179, 0.0421 1.074
C11H11Br0.18I0.82N4O 333.68 0.27 × 0.25 × 0.05 monoclinic P21/c 12.1261(5) 7.7747(3) 13.5613(6) 90 102.7210(10) 90 1247.13(9) 4 173 1.777 56.58 9926 3098 155 0.0184, 0.0489 1.055
C11H11I0.78N4.22O1.66 327.86 0.47 × 0.31 × 0.12 monoclinic P21/c 12.2248(10) 7.6978(6) 13.6339(11) 90 102.8190(10) 90 1251.03(17) 4 173 1.741 56.64 9706 3108 161 0.0342, 0.0897 1.069
C11H11Cl0.39I0.61N4O2.56 331.43 0.57 × 0.31 × 0.08 monoclinic P21/c 12.3212(17) 7.8240(11) 13.727(2) 90 102.897 90 1289.9(3) 4 173 1.707 56.64 10036 3209 191 0.0383, 0.0980 1.075
C11H11Cl0.53I0.47N4O3.12 327.59 0.31 × 0.25 × 0.09 monoclinic P21/c 12.309(2) 7.7389(13) 13.703(2) 90 102.853(3) 90 1272.6(4) 4 173 1.710 56.66 9756 3167 171 0.0563, 0.1462 1.040
likely stems from the systematic underestimation of lattice energies by DFT (due to underestimation of dispersion interactions), but also from our gross approximations made in estimating both the theoretical and experimental ∆G0exch values. Despite their limited quantitative accuracy, such computational analyses provide useful insights into the factors governing anion selectivities in competitive crystallizations, through deconvolution of the overall separation processes into simple steps that can be more readily evaluated energetically. For the current system, the gained insight was that (i) the anion binding in 1 · HX is not sufficiently strong to compensate for the unfavorable anion dehydration energy, and therefore the prevalent solvation (Hofmeister) bias dominates the anion selectivity, and (ii) the high flexibility of the H-bonded framework, manifesting as intraand interhelical extensions and contractions to accommodate the various anions, precluded shape recognition and peak selectivity, as indicated by the calculated lattice energies from the gas phase that increased monotonically with the anions’ charge densities. Conclusions This paper described an example of anion separation with H-bonded frameworks specifically functionalized with H-
bonding groups for anion binding. The frameworks were designed so that, through hierarchical H-bonding, the framework assembly and anion binding occur orthogonally and without interference. The formation of an isomorphous series of hydrogen-bonded helices including a wide variety of anions of different sizes, shapes, and basicities, offered a rare opportunity to understand anion selectivity in such systems, through competitive crystallization experiments and analysis of the resulting mixed crystal by single-crystal X-ray diffraction. Complementary theoretical calculations provided critical information about the energetics of anion binding and selectivity, which were found to be dominated by anion solvation manifested in the form of the Hofmeister bias. Although the solvation bias was significantly attenuated by the relatively strong anion binding in the crystals through urea H-bonds, CH · · · Xhydrogen bonds, and anion interactions with the pyridinium rings (X- · · · C bonds), it was not completely overcome. Another apparent factor preventing peak selectivity and shape recognition in these H-bonded crystals was their high flexibility manifested through helical extensions and contractions and interhelical metric adjustments to accommodate the different sizes and shapes of the anions. While such flexibility is likely important for retaining the same crystal structure across the whole anion
1914 Crystal Growth & Design, Vol. 8, No. 6, 2008
series, it proved detrimental to anion selectivity, as it allowed the host framework to adjust its structure to optimize the binding of all competing anions. These results stand in direct contrast to our previous work on rigid H-bonded framework hosts, which showed peak anion selectivity and shape recognition as a result of their almost perfect complementarity and rigidity.5a Experimental Section All reagents and solvents were purchased from commercial sources and used without further purification. Ligand 1 was synthesized according to a previously described procedure.5e Anions, except iodide, were analyzed by ion chromatography using an IonPac AS16 analytical column (for ClO4-) and an IonPac AS-14 analytical column (for NO3-, SO42-, and Cl-) (Dionex, Sunnyvale, CA).17 Iodide analysis was performed by Desert Analytics, Tucson, Arizona. Safety Note. Perchlorate salts are potentially explosive. Although no problem was encountered in the present study, care must be exercised when handling such compounds. X-Ray Crystallography. Single crystals of 1 · HX (X ) Cl-, Br-, I , NO3-, ClO4-) were grown by dissolving 0.1 mmol of 1 in 2 mL of hot methanol (water for X ) ClO4-) and adding 0.1 mmol of HX (0.1 mL of 1 M aqueous solution). Plate-shaped crystals appeared upon cooling to room temperature and/or slow evaporation. Single crystals of the mixed 1 · H(I)y(X)1-y salts (X ) Cl-, Br-, NO3-, ClO4-) were grown by slow cooling of hot aqueous solutions containing 0.1 mmol of 1, 0.1 mmol of HX, and 0.1 mmol of NaI in 2 mL of water. Single-crystal X-ray data were collected on a Bruker SMART APEX CCD diffractometer with fine-focus Mo KR radiation (λ ) 0.71073 Å), operated at 50 kV and 30 mA. All structures were solved by direct methods and refined on F2 using the SHELXTL software package.18 Absorption corrections were applied using SADABS, part of the SHELXTL package. All non-hydrogen atoms were refined anisotropically. Hydrogen atom positions were determined based on the maximum electron density and using idealized C-H and N-H bond lengths, then refined isotropically with a riding model. Crystallographic data are listed in Tables 4 and 5. Competitive Crystallization. To a solution containing 1 (0.021 g, 0.1 mmol), 0.5 M aqueous H2SO4 (0.1 mL, 0.05 mmol), and 0.5 M aqueous HNO3 (0.1 mL, 0.05 mmol) in 1.5 mL of deionized water was added an aqueous mixture of NaCl (0.1 mL, 1 M, 0.1 mmol), NaBr (0.1 mL, 1 M, 0.1 mmol), NaI (0.1 mL, 1 M, 0.1 mmol), NaClO4 (0.1 mL, 1 M, 0.1 mmol), and NaNO3 (0.1 mL, 0.5M, 0.05 mmol). Plate-shaped crystals appeared within 1 h. A few single-crystal specimens were collected after 1 h, 24 h, and 10 days and analyzed by X-ray diffraction, which found virtually no variation in crystal compositions for all specimen studied, indicating phase homogeneity and thermodynamic equilibration. The bulk of the crystallized solid was filtered after 10 days and washed with water and methanol. Yield 0.016 g (49%). A sample of 8 mg was dissolved in 25 mL of deionized water and analyzed by ion chromatography for ClO4-, NO3-, Br-, Cl-, and SO42-, which found concentrations of 26, 7.9, 2.9, 0, and 0 mg/L, respectively. These values correspond to 0.26, 0.127, 0.037, 0, and 0 mmol/L, respectively. The remaining sample was subjected to iodide analysis by Desert Analytics, which found 0.52 mmol/L I-. Pairwise competition experiments: Aliquots of 0.1 mL solution of NaI (1 M) were added over aliquots of 2 mL hot aqueous solutions containing 1 (0.1 mmol) and 0.1 mmol of the corresponding HX acid (X ) Cl-, Br-, NO3-, ClO4-). Multiple crystals from each experiment were collected at different intervals of time (3 h to 65 days) and analyzed by single-crystal X-ray diffraction, which found virtually no variation in the crystals’ composition within each batch, indicating bulk phase homogeneity and thermodynamic equilibration. The precipitated solids were filtered after 2 months and washed with MeOH. The iodide separation factors (R) were estimated based on the yields of crystallizations and anionic compositions of isolated crystals (Table 3), as determined by single-crystal X-ray diffraction (vide supra). This method only allows for approximate values of R (and thereby ∆G0exch (aq)) to be obtained, as the exact yields of crystallizations are difficult to measure accurately due to the loss of solid upon washing. Lattice Energy Calculations. The Vienna Ab Initio Simulation Package19 was employed to perform density functional theory (DFT) calculations with plane wave bases and periodic boundary conditions, in order to obtain lattice energies. The PBE functional form of the
Custelcean et al. generalized-gradient approximation was used for electron exchange and correlation.20 Projector-augmented wave method was used within the frozen core approximation to describe the electron-core interaction.21 The kinetic energy cutoff for plane waves was set at 450 eV. Only the Γ-point was used to sample the Brillouin zone of the lattice structures. The cell parameters were fixed at the experimental values, while all atoms in the cell were allowed to relax within the constraint of the experimental crystal symmetry. The convergence criterion for geometry optimization was such that the maximum force on an atom was less than 0.025 eV/Å. The energy of an isolated ion (cation or anion) was calculated by placing the ion in a large cubic box (for example, 25 × 25 × 25 Å3), and corrections for a charged system in a periodic box were applied so that the resultant energy approximate that of a truly isolated ion.
Acknowledgment. This research was sponsored by the Division of Chemical Sciences, Geosciences, and Biosciences, Office of Basic Energy Sciences, U.S. Department of Energy, under contract number DE-AC05-00OR22725 with Oak Ridge National Laboratory managed by UT-Battelle, LLC. Supporting Information Available: Cartesian coordinates for fragments used to obtain B3LYP/6-311+G** interaction energies for the six Cl- binding motifs (I-VI), and X-ray crystallographic information files (CIF). This information is available free of charge via the Internet at http://pubs.acs.org.
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