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Genevieve M. White
Comprehensive Paper
The Catholic University 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
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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 noncoordinating anions present, and solvent occlusion/hydrogen bonding, temperature, and
pressure. Classic examples from the literature, including Fe(phen)2(NCS)2 and bis(2,6bis(pyrazol-3-yl)-pyridine)iron(II) nitroprusside will be used as examples of spincrossover 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-2aminoethylamine (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
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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.
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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 spincrossover 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
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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.
N
NCS
N
Fe(II)
NCS
N
N
Figure 2 - Structural diagram Fe(phen)2(NCS)2 complex (3)
Spin state
LS
LS
LS
HS-2
HS-1
Experimental Conditions
30 K
130 K
1.0 GPa
30 K
293 K
Fe-N1 (Å)
1.990
2.014
1.975
2.177
2.199
Fe-N2 (Å)
2.007
2.005
2.003
2.184
2.213
Fe-N3 (Å)
1.953
1.958
1954
2.006
2.057
Table 1 - Bond distances (HS and LS) for Fe(II) complex Fe(phen)2(NCS)2 under 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 LS 1A1g, 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 hightemperature (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)
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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)2 are shown in Table 2. Also included are typical values for the
[Fe(phen)2]2+ cation for comparison.
Isomer shift δ [mm/s]
Quad. Splitting ΔEQ [mm/s]
HS [Fe(phen)3]2+
LS [Fe(phen)3]2+
HS Fe(phen)2(NCS)2
LS Fe(phen)2(NCS)2
0.8-1.4
1.7-3.1
0.45
0.30
0.98
2.67
0.37
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 of non-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 N4O2 and investigating the effects of this change.
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Such work has been done, for instance, with iron (II) triflate and TPA (tris-(2pyridylmethyl)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)
N
O
CD3
N
N
Fe(II)
N
N
N
N
CDCl3
N
CD3CN
CD3
N
O
Fe(II)
O
N
S
O
O
S
O
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
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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 noncoordinating 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.
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N
N
NN
Fe(II)
N
H
Fe(II)
N
N
N
H
O
N
N N
H
N
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(3aminopropyl)amine (trpn), tris(2-aminoethyl)methane (TRAM) or tris(2aminoethyl)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
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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.
CH3
+
NH2
N
3
Cl
TAMACl . 3HCl (H3L’)
N
NH2
NH2
HC
3
TRAM
N
3
TREN (H3L)
NH2
3
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.
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2+
N
N
Fe
NH
N
N
NH
N
N
NH
N
Figure 7 - Line Drawing of tripodal FeTren complex cation
2+
H
3+
+ CH3
N
N
N
N
N
NH
NH
N
N
NH
N
Fe
N
Fe
N
N
N
NH
N
NH
A
NH
B
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.
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H
O
NH
N
H
O
H
O
N
N
N
H
N
NH
A
B
O
NH
D
C
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.
O
H
HN +
NH
X
O
O
H
Cl
O
NH
O
H
NH
N
C
H
N
HO
N
O
M(II)
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
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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 Diego CA. Tris(2aminoethyl)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 57CoRh 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
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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-
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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
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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.
Synthesis of [MnH3L]K(ClO4)3: 2-imidazolecarboxaldehyde (0.385 g, 4.010 mmol) and
tris(2-aminoethyl)amine (0.2mL, 1.338 mmol) were refluxed in 55 mL methanol. After
10 minutes, the solution was a bright clear yellow. This was separated into three 10 mL
portions. MnCl2 ∙ 4H2O (0.090 g, 0.5003 mmol) in 5-10mL methanol was added to one of
these. There was no color change. KCl (0.100 g, 1.341 mmol) solid was added to the hot
solution. The solution was refluxed, and solid KCl remained on the bottom of the flask.
NaClO4 (0.790 g, 5.643 mmol) was dissolved in 30 mL methanol. One-third of this
solution (10mL, 1.881 mmol) was added. The reaction mixture was swirled and set aside
in the hood. A white solid began to precipitate.
Results and Discussion
This research focused on varying both 1) the Schiff-base elements (the backbone and
aldehyde); and 2) the double-salt elements (the central cation and the metal of the
imidazole complex) to see if the unique hydrogen-bonding feature described above could
be reproduced. Whether these structures were observed for complexes formed from tren
depended partly on the metal used to complex the ligand. With Fe(II), several H-bonded
double salts were prepared, using tren and 2-imidizolecarboxaldehyde. A HS Mn(II)
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simple complex was also isolated, but did not exhibit the hydrogen-bonding feature
sought. Instead, the complex hydrogen-bound in a monodentate fashion, and was not a
double salt as were the other complexes. Both the iron (II) complexes of tren and TRAM,
when condensed with 2-imidazolecarboxaldehyde, gave double salts with KClO4.
Subsequent discussion will focus on the tren complex rather than the TRAM. The second
preparation of a double salt with the TRAM backbone is mentioned to demonstrate that
the formation reaction is general.
Analogously to complexes described in the introduction, the spin state of the
originally HS Fe(II) tren compex investigated in this research is effected by donor atoms,
ligand flexibility, and solvent occlusion. Just as in the example of solvent occlusion
above, in the tren complex the perchlorate hydrogen bonds to the imine hydrogen and the
imidazole hydrogen. The perchlorate then pulls away the imidazole hydrogen and imine
hydrogen, making the ligand a stronger-field ligand and locking the complex into the LS
state (as in Figure 5). This lock into LS conformation may partially depend on the size of
the counterion, which in successful experiments was perchlorate. Finally, it was found
that the atomic radii of the counterion to perchlorate is also key in determining whether
the complex would form or not.
Mössbauer
Mössbauer results – isomer shifts δ and quadropole splitting ΔEQ - for
[FeH3L](ClO4)2.H2O, [FeH3L]K(ClO4)3, [FeH3L]Rb(ClO4)3, [FeH3L]Cs(ClO4)3,
and[FeH3L]NH4(ClO4)3 taken at room temperature are shown in Table 3 below. Isomers
shifts at 0.29 mm/s and quadropole splittings at ~0.35 mm/s are indicative of a LS
complex when compared with the HS and LS values for the spin crossover complex
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[FeH3L](ClO4)2.H2O, and are also consistent with predictions made concerning the
change in field strength promoted by partial deprotonation of the imidizole and imine
nitrogens due to H-bonding with perchlorate as was demonstrated in Figure 5 above.
[FeH3L](ClO4)2.H2O
[FeH3L]K(ClO4)3
[FeH3L]Rb(ClO4)3
[FeH3L]Cs(ClO4)3 *
[FeH3L]NH4(ClO4)3
δ [mm/s]
0.21 (LS)
1.00 (HS)
0.29
0.29
0.29
0.29
ΔEQ [mm/s]
0.39 (LS)
2.22 (HS)
0.37
0.36
0.35
0.35
Table 3 - Values for HS/LS Fe(II) double salts. *Spectrum was noisy due to small sample size.
All Mössbauer spectra showed the single doublet which is consistent with the
presence of LS Fe(II), in contrast to the spectra taken of the parent complex
[FeH3L](ClO4)2.H2O alone, which showed doublets corresponding to both HS and LS
Fe(II). Spectra for the [FeH3L]K(ClO4)3 and [FeH3L](ClO4)2.H2O are shown in Figure 11
below. This indicates that formation of the double salt shifts the position of the
LS(1A)↔HS(5T) equilibrium of the original spin-crossover complex from 62.8% HS to
entirely LS.
Figure 11 – Lorentzian site analysis of [FeH3L]K(ClO4)3 (right) and [FeH3L](ClO4)2.H2O (left)
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X-Ray
All structures exhibited a supramolecular symmetry in the space group Pbar3. Xray diffraction data for [FeH3L]K(ClO4)3 produced the ORTEP diagram shown in Figure
12 of the iron complex cation. The central metal is the LS form of Fe(II). ORTEP
diagrams for the other complexes are similar and so will not be provided.
Figure 12- ORTEP diagram of [FeH3L]2+
As was mentioned in the introduction, there are structural clues as to the spin state
of the metal of the imidazole complex (in this case, Fe(II)): M-Nimidazole and M-Nimine
bond distances, the Nimidazole-M-Nimine bite angle, the Nimidazole-M-Nimine’ trans angle, and
the M-Nap distances. This crystallographic data (taken at for the new complexes is listed
in Table 4 below, shows a shortening of M-Nimidazole and M-Nimine bond distances, an
increase in the bite angle, an increase in the trans angle, and lengthening of the M-Nap
non-bonding distances, all of which indicate a LS Fe(II) center. These are offered for
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comparison with the data for the original spin-crossover complex, which is pure LS at
this temperature (100 K).
M-Nimidazole
(average)
[FeH3L]K(ClO4)3
1.9432
[FeH3L]Rb(ClO4)3
1.9552
[FeH3L]Cs(ClO4)3
1.957
[FeH3L]NH4(ClO4)3 1.9578
[FeH3L](ClO4)2.H2O 1.964
M-Nimine
(average)
1.9624
1.9740
1.978
1.9751
1.980
Bite angle
81.08
81.03
81.08
81.08
80.77
Trans
angle
173.02
172.82
172.80
172.91
173.19
M-Nap
3.350
3.376
3.386
3.380
3.436
Table 4 - Crystallographic data for new complexes
Further analysis of the structure revealed more similarities between the
[FeH3L]K(ClO4)3,[FeH3L]Rb(ClO4)3, [FeH3L]Cs(ClO4)3, and [FeH3L]NH4(ClO4)3
complexes. The [FeH3L]K(ClO4)3 will serve as an example to demonstrate this structural
feature. In this complex, each potassium ion is coordinated to six different perchlorate
anions, each acting as a bidentate donor to the ion, forming a distorted icosahedral cage
repeated throughout the lattice. (Figures 13, 14) Formation of such cages is not
unprecedented. (12)
Figure 13 The hydrogen bonding arrangement between one perchlorate anion and three [FeH3L]2+
complexes (left) and one [FeH3L]2+ complex and three perchlorate anions (right).
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Figure 14- Coordination of the potassium ion by the six perchlorate anions. Top left shows the
octahedral arrangement of the chlorine atoms around the potassium. Top right shows the CN=12
potassium ion. Bottom left draws in the oxygen-oxygen non-bonded distances to show the framework.
Bottom right is the same as bottom left but eliminates the chlorine and oxygen atoms not bound to
potassium.
Furthermore, the overall structure contains similarities between the double salts. An
C3i (S6) axis forms the symmetry axis through the cell. Two independent cations are
arranged on this axis, and midway between the cation positions are three symmetryrelated perchlorate anions. The three chlorine atoms form an equilateral triangle
perpendicular to the c axis (Figure 15). This pattern extends ad infinitum, forming what
could be called a 1-dimensional polymer.
X
X
X
O
O
X
M
X
M
X
M
O
X
X
X
O
X
B
A
Figure 15 - (A) 1-D polymer, where M=Fe(II) and X=Cl; (B) detail of perchlorate, where X=Cl
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Overall, this polymer has a general formula of the form [M(ClO4)32-]X, which requires
a cation of 2+ charge such as Fe(II) or Mn(II) to stabilized the “anion”. Both these 2+
metals were attempted, but only Fe(II) gave the desired double salt. Mn(II) failed to give
the double salt under analogous reaction conditions. This failure could possibly be
attributed to either differences in solubility between Fe(II) and Mn(II), the geometry of
the bidentate hydrogen-bonding donor site, the complex acidity, or electronic differences
between the two 2+ metals.
It was also found that the formation of this structure is heavily dependent on the
atomic radii of the central cation. Na+, K+, Cs+, Rb+, and Ag+ were used in order to vary
the size of the perchlorate counterion. Atomic radii are shown in Table 5 below.
Structures incorporating K+, NH4+, Rb+, and Cs+ - which have atomic radii between 1.78
and 2.02 Å - were obtained, with the ions fitted into similar distored icosahedral cages.
Radii (Å)
K+
1.78
Rb+
1.86
Cs+
2.02
NH4+
1.51Φ
Na+
1.53*
Ag+
1.42Γ
Table 5 - Atomic radii of central cations. All values are for CN=12 unless otherwise noted. ΦDoes not
take into account hydrogen bonding, which would increase the radius. *The value for NH4+ is the
thermochemical radius, calculated using a Born-Haber calculation and enthalpies of formation. Γ
The radius is for CN=8. This radius would be greater for CN=12 as the ionic radius increases with
increasing coordination number. (13)
No double salt complexes were obtained using Na+ or Ag+ ion. It is supposed that Na+
is too small to form the icosahedral cage formed by perchlorate (Figure 16). This
indicates that the cages’ ability to expand or contract has limits: even if the cage were
able to form, the bond angles would be highly distorted and the complex would simply
fall apart due to strong repulsive forces between perchlorate anions. Expansion of the
cage to accommodate larger cations (Cs+ and Rb+) is more feasible than contraction of
the cage for these reasons. Another reason for the failure of Na+ to precipitate in the
double salts could be the solubility of sodium perchlorate in methanol relative to other
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alkali metals. However, the successful formation of a double salt of NH4+, whose
solubility in methanol is higher than that for Na+, suggests that atomic radii is the
determining factor of whether or not the double salt will form.
In the case of Ag+(which is presumably similar in size to K+, Rb+, and Cs+ at
CN=12 because bond length increases with increasing coordination number as a result of
crowding), the cation proved too strong an oxidizing agent and presumably oxidized the
Fe(II) complex to Fe(III) instead of forming a similar complex to the others in the series.
The reaction of the Schiff-base with silver perchlorate in methanol gave an immediate
color change from red to green-dark blue. The precipitated blue powder was not further
characterized, but it seems that redox properties of the central metal must be taken into
consideration in addition to size in the formation of these double salts.
Figure 16 - Icosahedral "cage" formed by perchlorates
Finally, one valuable aspect of this investigation has been the development of a
new perchlorate anion binding site. In the design of binding sites, a host-guest mentality
usually prevails, with a host being designed to fit the anion guest. (14, 15) Providing
perchlorate anions with hosts is a particularly difficult task for several reasons including
the problem of providing ample hydrogen bonding sites between host and guest. One
example that shows the intricacy of these binders is illustrated in Figure 17. These tren
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and TRAM complexes prepared in this research are a new example of host-guest
perchlorate anion binding, for they provide both hydrogen bonding sites to the 2imidazolecarboxaldehyde anion-binder as well as space to attach itself.
Figure 17 - Synthesized ClO4- binder 1,4,5,7,16,19,22-hexamethyl-1,4,7,16,19,22hexaaza[9.9]paracylophane (left) and its precursor (right), exemplifying the intricacy required for
the design of such anion binders. (15)
Conclusion
In this research, various components of a particular series of double-salts were
varied to investigate the influences of these components on the self-assembly of the
complexes, particularly their unique structural features. These unique features include a
non-conventional hydrogen bonding pattern, the formation of caged ions (K+, Cs+, Rb+,
and NH4+), the formation of the larger structure into a one-dimensional polymer repeating
ad infinitum, and the function of the complex as a perchlorate anion binder. Five
complexes exhibiting these structural features, [FeH3L](ClO4)2.H2O, [FeH3L]K(ClO4)3,
[FeH3L]Rb(ClO4)3, [FeH3L]Cs(ClO4)3, and [FeH3L]NH4(ClO4)3 were synthesized. After
experiment, it was found that all of the unique features mentioned above depend largely
on two key properties of the caged central cation: the size of the cation as well as its
oxidizing strength.
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Acknowledgements
I would like to thank the following researchers for their invaluable contribution to this
work:
Gregory Brewer, Ph.D., CUA (Research Advisor)
Cynthia Brewer, Ph.D., CUA (Research Advisor)
Carol Viragh, Ph.D., VSL (Mössbauer measurements)
Ray Butcher, Ph.D., Howard University (Crystallographic measurements)
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