WHITE G
Genevieve M. White
Comprehensive Paper
The CatholicUniversity of America
Washington, D.C.
Hydrogen Bonding in Supramolecular Complexes: Supramolecular assemblies prepared from an iron(II) tripodal complex containing 2-imidazolecarboxaldehyde
April 4, 2008
Abstract
The spin crossover phenomena (SCO) is one that is exhibited by the first-row
d4-d7 transition metals. These metals can have two different electronic ground states that are so close in energy that only a small change is necessary to switch electronic states. The high spin state contains the highest number of unpaired electrons, and the low spin state contains the lowest number of unpaired electrons. The spin state of the complex can be changed from one electronic state to another by varying different components of the complex, including the donor atoms of the ligands surrounding the metal, the non-coordinating anions present, and solvent occlusion/hydrogen bonding, temperature, and pressure. Classic examples from the literature, including Fe(phen)2(NCS)2 and bis(2,6-bis(pyrazol-3-yl)-pyridine)iron(II) nitroprusside will be used as examples of spin-crossover systems.
In this research, the formation of double salts was attempted by the reaction of a Fe(II) Schiff-base complex with various metal perchlorates:
[FeIIH3L](ClO4)2 + MClO4 → [FeIIH3L]M(ClO4)3
where M = K+, Rb+, Cs+, NH4+, Ag+, Na+ and H3L is the Schiff-base condensate of tris-2-aminoethylamine (tren) with 2-imidazolecarboxaldehyde. The parent iron complex at room temperature is mostly HS, but on formation of the double salts becomes entirely LS at room temperature. These double salts exhibited several unusual and interconnected structural features. These features include: (1) a unique bidentate hydrogen-bonding pattern between perchlorate and the organic ligand; (2) distorted icosahedral cages around the alkali cation; (3) the formation of what could be called a one-dimensional polymer (MX32-). The Mössbauer and bond distances of all five complexes indicate a low-spin state for Fe(II). It was found that certain alkali cations have the ability to arrange perchlorates and iron complexes into a supramolecular assembly, and that the formation of such assemblies is reliant on the atomic radii of the alkali cations.
Introduction
The spin-crossover phenomenon (SCO) is a research topic that has recently gained new importance in the fields of inorganic and coordination chemistry. There are two spin states for the metal, high spin (HS) and low spin (LS), and the spin state of a metal in a particular complex depends on the magnitude of the energy gap between the t2g (bonding) and eg (antibonding) orbitals, ∆0 (Figure 1). If ∆0 is greater than the interelectronic pairing energy, P, the metal becomes low spin, and conversely if ∆0 is less than P, the metal becomes high spin. (1) Some complexes exhibit behavior somewhere between HS and LS and these are known as “spin-crossover” complexes. It has been found that these complexes can be switched between spin states by thermal, pressure, and even light changes. (2) This allows for the potential use of spin-crossover complexes in a broad range of applications, from switches to memory devices. (2) However, the spin-crossover phenomenon is only observed in the first row d4-d7 octahedral metals, with Fe(II) and Fe(III) among some of the most commonly used in experiment. Because it is diamagnetic (as shown in the figure below), Fe(II) exhibits some of the most profound effects of spin crossover from the paramagnetic to the diamagnetic state.
Figure 1 - Orbital Diagrams showing HS (left) and LS (right) states for a d6 metal
In such spin crossover complexes, there are several ways to determine the spin state of the metal. Bond distances, magnetic moment, and Mossbauer spectra are examples of these methods. What follows are data for a well-known Fe(II) spin crossover complex, Fe(phen)2(NCS)2, shown in Figure 2 (3). Note that bond distances in general increase with the transition from LS to HS. This would be due to the movement of two electrons from the t2g to the eg orbital, giving a increase in ionic radius of the complex. Bond distances for the Fe(phen)2(NCS)2 complex are given in Table 1, where the dramatic change in bond lengths can be clearly seen.
Figure 2 - Structural diagram Fe(phen)2(NCS)2 complex (3)
Spin state / Experimental Conditions / Fe-N1 (Å) / Fe-N2 (Å) / Fe-N3 (Å)LS / 30 K / 1.990 / 2.007 / 1.953
LS / 130 K / 2.014 / 2.005 / 1.958
LS / 1.0 GPa / 1.975 / 2.003 / 1954
HS-2 / 30 K / 2.177 / 2.184 / 2.006
HS-1 / 293 K / 2.199 / 2.213 / 2.057
Table 1 - Bond distances (HS and LS) for Fe(II) complex Fe(phen)2(NCS)2under various indicated experimental conditions. (HS-2 indicates a thermal HS state, while HS-1 indicates a light-induced metastable HS state.) (4)
Second, a typical spin-crossover plot of magnetic moment (μeff) vs. temperature will look like that in Figure 3 for Fe(phen)2(NCS)2. At lower temperatures, the complex is LS1A1g, switching over to HS 5T2g at ~ 174K. (5) Also shown in the figure is the plot of magnetic moment vs. temperature for a similar complex, Fe(phen)2(NCSe)2. The high-temperature (HS) limit for the magnetic moment of the Fe(phen)2(NCS)2 complex is found at 5.20±0.05 BM at 430 K, and the low-temperature limit at 0.84±0.01 BM at 77.2 K. (6)
Figure 3 - Plot of magnetic moment vs. temperature for spin crossover Fe(II) complexes Fe(phen)2(NCS)2 (filled and empty circles) and Fe(phen)2(NCSe)2 (half-circles) (6)
Lastly, Mössbauer spectra are indicative of the spin state of the metal. Values for isomer shifts and quadropole splitting for HS and LS Fe(II) Mössbauer spectra of Fe(phen)2(NCS)2are shown in Table 2.Also included are typical values for the [Fe(phen)2]2+ cation for comparison.
HS [Fe(phen)3]2+ / LS [Fe(phen)3]2+ / HS Fe(phen)2(NCS)2 / LS Fe(phen)2(NCS)2Isomer shift δ [mm/s] / 0.8-1.4 / 0.45 / 0.98 / 0.37
Quad. Splitting ΔEQ [mm/s] / 1.7-3.1 / 0.30 / 2.67 / 0.34
Table 2 - Typical isomer shifts and quadropole splittings for Fe(II) and Fe(III). HS values were taken at 293 K, and LS values at 77 K. (6)
Several factors are involved in determining whether a spin-crossover complex is HS or LS, including the donor atoms of the ligands surrounding the metal (for example, C and N are both strong-field ligands and hence favor the LS state, while O and the halides are weak-field ligands and favor the HS state), the presence ofnon-coordinating anions, and solvent occlusion/hydrogen bondin.
One example of the influence of changing the donor set on spin state can be seen by forming a complex in the presence of (a) a strong-field donor set such as N6 as well and (b) a weak-field donor set such as N4O2and investigating the effects of this change. Such work has been done, for instance, with iron (II) triflate and TPA (tris-(2-pyridylmethyl)amine) complexes. (7) These two when mixed form a six coordinate product with two coordination sites open potentially for anions. It was discovered that, when dissolved in a solvent such as deuterated CD3CN with an N6 donor set, a LS complex formed (Figure 4, left). Furthermore, if this LS complex were dried and redissolved in a solvent such as CDCl3 or (CD3)2CO with an N4O2 donor set,a HS complex is formed (Figure 4, right). There is an equilibrium (spin-crossover) relationship between the two, as HS can also be re-converted to LS if it is redissolved in the first solvent. Seen here is the influence of the strength of the donor set: N is the strong-field ligand and the N6 donor set causes the complex to go LS, and O is a weak-field ligand and thus the N4O2 donor set influences the LS complex to switch spin-states to HS. This illustrates the influence of strong and weak-field ligands on the spin state of the central metal. (1)
Figure 4 – LS with N6 donor set (left) – HS with N4O2 donor set (right) to illustrate spin crossover dependence on strong and weak field ligands
The influence of the presence of non-coordinating anions on spin state is an area of coordination chemistry not fully understood, but there are illustrative examples found in the literature. Even when all coordination sites are filled and no further bonding is possible, the mere presence of a particular ion in the crystal lattice can affect the spin state of the complex. For example, one research group varied the non-coordinating anion in a double salt of a heptadentate complex and observed an influence of these anions on the spin state of the complex. Complexes of the formula [FeL][X]2.H2O were synthesized, where L=tris(4-{pyrazol-3-yl}-3-aza-3-butenyl)amine and X = BF4-, ClO4-, NO3- and CF3SO3-. Salts incorporating BF4- and ClO4- into the crystal lattice gave HS complexes between 5-300 K, while the other two were HS at room temperature but underwent a HS→LS spin transition upon cooling. The relationship between the presence of non-coordinating anions and spin state is least understood of the three factors affecting spin state listed above. (8)
Thirdly, instances have also been reported where hydrogen-bonding from solvent occlusion has heavily influenced the spin state of the central metal. For example, there are a number of hydrated salts that have been characterized where spin state switches between a singlet 1A1 and a quintet 5T2 between the anhydrated and hydrated forms of the salt. (9) Closer investigation revealed a dependence of spin state on solvent binding. Figure 5 illustrates an example the potential effects of solvent binding on spin state. The Fe(II) in this complex was HS, and pyrazole is a relatively weak-field ligand. But in hydrogen bonding to the solvent H2O, the hydrogen was pulled away from the pyrazole molecule, effectively making the pyrazole a stronger-field ligand and forcing the complex to go LS (∆0Im->∆0ImH). This same argument is also valid for imidazole.
Figure 5- Solvent occlusion and hydrogen bonding shown in bis(2,6-bis(pyrazol-3-yl)-pyridine)iron(II) nitroprusside. On the left is a drawing of the entire [Fe(3-bpp)2]2+ complex with hydrogens omitted for clarity, and on the right the hydrogen-bonding with the solvent is shown in detail. (9)
In Schiff-base reactions, there are two basic components - the amine and the aldehyde or ketone - which condense to form a stable imine with the general formula R1R2C=N-R3. Previous work done involving an investigation of the effects of variation of these two Schiff-base components on spin state includes the preparation of Schiff-base spin-crossover iron complexes where R = tris(2-aminoethyl)amine (tren), tris(3-aminopropyl)amine (trpn), tris(2-aminoethyl)methane (TRAM) or tris(2-aminoethyl)methyl ammonium chloride (TAMACl) (Figure 6)and with various aledehydes. These ligands were prepared by Schiff base condensation of three molar equivalents of an imidazole carboxaldehyde with either TRAM, TAMACl, trpn, or tren. Advances made in this research include the synthesis of a seven-coordinate TRAM complex. (10) The research presented in this paper also focused on the synthesis of the above-mentioned ligands via a Schiff-base condensation, but in particular aimed to investigate and prepare double salt complexes exhibiting unique and non-conventional bidentate hydrogen bonding patterns.
TAMACl . 3HCl (H3L’) TRAM
TREN(H3L)TRPN
Figure 6 – Backbone ligands used for Schiff-Base condensates
Tren was chosen as the backbone for these Schiff-base condensations based on its availability and ease of synthesis (Figure 7). TRAM, TAMACl, or trpn (Figure 8) could also have easily been chosen as the backbone. To demonstrate this feature, one complex was synthesized using TAMACl to see if the non-conventional hydrogen bonding pattern could be reproduced with a backbone other than tren. In fact, a complex was successfully synthesized and its crystal structure determined. It was found that it exhibited the same unique hydrogen-bonding feature found in the tren complexes.
Figure 7 - Line Drawing of tripodal FeTren complex cation
Figure 8 – Line drawings of the other possible FeTRAM (A) and FeTAMACl (B) complex cations
Furthermore, although several aldehydes could have been used in the Schiff-base condensation, only 2-imidazolecarboxaldehyde (Figure 9A) had the potential to hydrogen bond to both imine and imidazole nitrogens to reproduce the bidentate non-conventional hydrogen bonding feature. Neither B, C, or D were attempted because the arrangement of nitrogens around the ring would not allow for the hydrogen-bonding desired.
Figure 9 - Aldehyde component of Schiff-base
This positioning of the hydrogen-bonding nitrogens on the aldehyde is structurally similar to the guanidinium group found on the amino acid arginine (Figure 10). (11) This may have implications for future applications of this research in biochemistry and bioinorganic chemistry.
Figure 10 - Line drawings for comparison of Guanidinium group of Arg and H-bonding to
2-imidazolecarboxaldehyde
Aside from the Schiff-base component of these iron(II) tripodal complexes are the components of the resulting double salt which can also be varied in experiment. These include the central cation (the counterion to ClO4-) and the metal in the imidazole complex. In the course of this research, complexes using central cations Na+, K+, Cs+, Ag+, Rb+, and NH4+ were attempted in the course of research with varied results. For the metal in the imidazole complex, two M2+ metals (Fe2+ and Mn2+) were attempted. The choice of M2+ (rather than M+ or M3+, for instance) will be discussed later.
In the research presented in this paper, both ligands and counterions of double salts were varied to investigate the effects of change on the spin state of the central iron as well as on hydrogen bonding in the supramolecular structure.
Experimental
Elemental analyses were determined by Galbraith Laboratories, Knoxville, TN. Mass spectal analyses were obtained from HT Laboratories, San DiegoCA. Tris(2-aminoethyl)amine, 2-imidazolecarboxaldehyde, rubidium perchlorate, rubidium chloride, caesium perchlorate, caesium chloride, and ammonium perchlorate were obtained from Aldrich. Sodium perchlorate monohydrate, potassium chloride, and potassium perchlorate were obtained from Fisher. All solvents were of reagent grade and used without further purification.
The 57Fe Mössbauer spectra were recorded from powdered samples with a constant acceleration MS1200 Ranger Scientific spectrometer and a ca. 1.85 GBq 57Co-Rh source. The sample thickness was ~50-80 mg cm-2. The line width of the calibration spectrum was 0.29 mm s-1. The chemical isomer shift data are quoted relative to the centroid of the metallic iron spectrum at room temperature. The data were analyzed by a constrained least squares fit to Lorentzian shaped lines.
Crystal data for all complexes were collected on a Bruker Apex 2 diffractometer. All structures were solved using direct methods program SHELXS-97. All non-solvent heavy atoms were located using subsequent difference Fournier syntheses. The structures were refined against F2 with the program SHELXL, in which all data collected were used including negative intensities. All non-solvent heavy atoms were refined anisotropically. All hydrogen atoms were located by Fournier difference except for the ammonium hydrogen atoms at 293K. The hydrogen atom of the ammonium cation was located at 173 K.
Syntheses of characterized complexes
Synthesis of [FeH3L]K(ClO4)3: 2-imidizolecarboxaldehyde (0.386 g, 4.021 mmol) was added to 30 mL methanol. 0.20 mL tris(2-aminoethyl)amine (0.1954 g, 1.338 mmol) was added and the solution was refluxed. After 12 minutes, the solution was a bright clear yellow. FeCl2 ∙ 4H2O (0.267 g, 1.343 mmol) was added, and the solution immediately turned a dark red. After 10 minutes further refluxing, the solution was divided into three 11 mL portions. Two different preparations were performed with potassium. KCl (0.033 g, 0.4427 mmol) was added to each. Solid KCl still remained on the bottom of the flasks.
NaClO4(0.254 g, 1.814 mmol) dissolved in 5-10mL was added. The flask was swirled and set aside in the hood. One week later, 254 mg product was collected by filtration, and samples were sent off for Mossbauer and x-ray crystallography. Elemental analysis calculated for C18H24N10Cl3FeKO12: C 27.94, H 3.13, N 18.10. Found: C 27.40, H 2.94, N 17.59.
Synthesis of [FeH3L]Rb(ClO4)3: 2-imidazolecarboxaldehyde (0.402 g, 4.188 mmol) was dissolved in 75 mL methanol. Tris(2-aminoethyl)amine (0.2 mL, 1.397 mmol) was added and the solution was refluxed for 15 minutes. The solution turned a bright yellow. Methanol was added to bring the volume to 75mL and this was divided into three 25 mL portions. Fe(ClO4)2 ∙ 6H2O (0.169 g, 0.4656 mmol) dissolved in 5-10mL methanol was added to one portion of the hot solution. The solution immediately turned a dark blood-red. RbClO4(0.260 g, 1.405 mmol) was added in a slurry of methanol. This red solution was refluxed for 10 minutes and a white solid began to come out. 10 mL methanol was added with no effect. Elemental analysis calculated for C18H24N10Cl3FeRbO12: C 26.36, H 2.95, N 17.08. Found: C 26.48, H 3.07, N 17.03.
Synthesis of [FeH3L]Cs(ClO4)3: Tren (1.005 g, 6.884 mmol) was weighed out and transferred to a larger beaker with 80 mL methanol. 2-Imidazolecarboxaldehyde (1.977 g, 20.594 mmol) was added and the solution was set up to reflux. An additional 80 mL of methanol was added to dissolve the solid. After fifteen minutes, the solution had turned a dark orange color. The solution was diluted to 200mL and divided into 3 portions, one of 120 mL and two of 40 mL. To one 40 mL portion Fe(ClO4)2. 6H2O (0.495 g, 1.365 mmol) was added, dissolved in methanol. The solution turned dark red. This solution was divided into two 35 mL portions, and to one portion was added solid CsClO4 (0.478 g, 2.057 mmol). This solution was set aside in the hood. Elemental analysis calculated for C18H24N10Cl3FeCsO12: C 24.92, H 2.79, N 16.14. Found: C 25.25, H 2.97, N 16.02.
Synthesis of [FeH3L]NH4(ClO4)3: Ammonium perchlorate (0.074 g, 0.631 mmol) was added as a solid to a refluxing solution of 1.H2O (0.200 g, 0.315 mmol) in methanol (50 mL). The dark red solution was set aside to concentrate. After 2 d, dark red crystals (0.128 g, 54%) suitable for X-ray diffraction were removed by suction filtration. Elemental analysis calculated for C18H28N11Cl3FeO12: C 28.72, H 3.75, N 20.47. Found: C 28.70, H 3.55, N 20.31.
Synthesis of [FeH3L](ClO4)2.H2O: TAMACl (0.050 g, 0.1634 mmol) was heated and stirred in 22 mL methanol. The solid did not dissolve. After refluxing for 5 minutes, 0.1M methanolic KOH (4.90 mL, 0.490 mmol) was added. After 45 minutes of refluxing and 50 mL methanol and adding 50mL more methanol, the solid had dissolved and the solution was pale yellow. 2-imidazolecarboxaldehyde (0.047 g, 0.4896 mmol) dissolved in 22 mL methanol was added and the solution was refluxed for 10 minutes. The solution was a clear, pale pink color. FeCl2 ∙ 4H2O (0.034 g, 0.1710 mmol) dissolved in 5-10mL methanol was added. The solution immediately turned a dark red-purple color. After 5 minutes refluxing, NaClO4(0.094 g, 0.6714 mmol) dissolved in methanol was added. Solution was swirled and set aside in the hood.