Synthesis and characterization of iron trisphenolate complexes with hydrogen-bonding cavities

Article abstract of DOI:10.1021/ic5002286

A new family of C₃-symmetric ligands, featuring phenolate donors and a secondary coordination sphere, have been synthesized. We report the synthesis and subsequent coordination chemistry of these new tripodal N-anchored tris(phenolate) chelates, [tris(5-tert-butyl-3-N-carboxamide-2-hydroxybenzyl) amines] (H₃^RSalAmi), to iron(II), iron(III), and zinc(II). These electron-rich complexes have intramolecular hydrogen bonds, and therefore the potential to stabilize biologically relevant substrates in small-molecule activation chemistry.

Full text of DOI:10.1021/ic5002286

Communication
pubs.acs.org/IC
Synthesis and Characterization of Iron Trisphenolate Complexes with
Hydrogen-Bonding Cavities
Mario Adelhardt, Matthew J. Chalkley, Frank W. Heinemann, Jorg Sutter, Andreas Scheurer,
̈
and Karsten Meyer*
Department of Chemistry and Pharmacy, Inorganic Chemistry, Friedrich-Alexander University Erlangen−Nurnberg (FAU),
̈
Egerlandstraße 1, 91058 Erlangen, Germany
S
* Supporting Information
shown that hydrogen bonding reduces the M−E bond order and
ABSTRACT: A new family of C3-symmetric ligands,
featuring phenolate donors and a secondary coordination
sphere, have been synthesized. We report the synthesis
and subsequent coordination chemistry of these new
tripodal N-anchored tris(phenolate) chelates, [tris(5-tert-
butyl-3-N-carboxamide-2-hydroxybenzyl)amines]
stabilizes high-spin species, making these species more similar to
what is spectroscopically observed in metalloproteins.16
Our group has recently worked to develop a series of C3-
symmetric N-anchored tripodal ligands with donors from
tris(carbenes) to tris(phenolates), including mixed-donor
species, for first-row transition metals.17 Subsequently, we
believed there was an opportunity to combine the biologically
relevant phenolate system18,19 with a biologically relevant
secondary coordination sphere. To this end, we developed a
new family of C3-symmetric N-anchored tris(phenolate) ligands
with amide moieties in the secondary coordination sphere
R
(H3 SalAmi), to iron(II), iron(III), and zinc(II). These
electron-rich complexes have intramolecular hydrogen
bonds, and therefore the potential to stabilize biologically
relevant substrates in small-molecule activation chemistry.
R
(H3 SalAmi). The synthetic pathway (see the Supporting
Information, SI) allows for a large-scale synthesis from
inexpensive precursors with convenient control of the sterics of
the binding cavity in the last step by changing the primary amine
used. The ligands are selectively deprotonated at the phenol
position by reaction with NaOMe and, as seen in Scheme 1,
he use of C3-symmetric ligands has become more
Twidespread because ligands of this type have repeatedly
shown the ability to stabilize fleeting intermediates and
previously inaccessible species.1 It has been shown that these
ligands, by lowering the symmetry around the metal from
octahedral or tetrahedral to trigonal, make additional non-
bonding orbitals available that, in turn, stabilize these highly
reactive terminal ligands.2 On iron in particular, C3-symmetric
ligands have provided access to both terminal nitride1,3−5 and
oxo ligands.6−8
Scheme 1. Synthesis of Metal Complexes from the
Deprotonated SalAmi Ligand and Metal Salts
Much of this work has been pursued because such species are
believed to be models of important intermediates in biological
transformations. In biological systems, these transformations are
almost always made more energetically favorable by the
utilization of hydrogen-bonding interactions between the
enzyme and substrate or intermediates. The hydrogen bonds
provide both a means of stabilization and a means of directing the
reactivity. Accordingly, there has been a growing interest in
developing ligands that incorporate hydrogen-bond donors into
the secondary coordination sphere in a biomimetic fashion. This
attempt to synthetically mimic nature has led to some impressive
successes, such as Tolman et al.’s copper hydroxide,9 Masuda et
al.’s copper hydroperoxo,10 Nocera et al.’s iron hydroxide,11
Berreau et al.’s zinc alkoxide,12 and Goldberg et al.’s iron
peroxide.13
Borovik et al. have applied both of these approaches through
their use of tris[(N′-tert-butylureaylato)-N-ethyl)]aminato
(H6buea), a C3-symmetric ligand with a secondary coordination
sphere.14 By using ligands with varying numbers of hydrogen-
bond donors from 0 to 3, Borovik et al. have demonstrated how
the secondary coordination sphere is essential to stabilizing a
variety of terminal ligands of biological interest.15 Furthermore,
his studies of the electronic structure of these complexes have
subsequently metalated with both iron(II) and iron(III) chloride
or iron(III) triflate salts to generate the monoanionic and neutral
complexes, respectively. Additionally, as expected, the iron(III)
species could be generated through chemical oxidation of the
iron(II) species. Tris(phenolate) amine complexes of iron are
relatively rare,17,20−23 and to this point, none has introduced a
secondary coordination sphere. For structural, spectroscopic,
and electrochemical comparison, a zinc(II) species was also
synthesized.
The divalent iron complex TMA[(tBuSalAmi)FeII] (1a; TMA
= tetramethylammonium) was fully characterized by means of
1
single-crystal X-ray crystallography, H NMR, Moßbauer, IR,
̈
UV−vis spectroscopy, and variable-temperature (VT)/variable
Received: January 29, 2014
Published: March 3, 2014
© 2014 American Chemical Society
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Inorganic Chemistry
Communication
field (VF) SQUID magnetization. The isopropyl derivative
TMA[(iPrSalAmi)FeII] (1b) was characterized by analogous 1H
NMR and IR. The molecular and crystal structure of 1a (see
Figure 1) was determined by X-ray crystallography. The
Figure 2. Zero-field Moßbauer spectrum of 57Fe-enriched 1a in solution
̈
at 77 K [top; fit results of δ = 1.19(1) mm/s and ΔEQ = 2.84(1) mm/s].
VT/VT SQUID data for 1a in the solid state (bottom; fit results of μeff,RT
= 5.1 μB) and VF SQUID data (inset: H = 1, 3, and 5 T; fit results of D =
6.472 cm−1 and E/D = 0.282).
Figure 1. Molecular structure of the complex anions of 1a in crystals of
TMA[(tBuSalAmi)FeII(pyridine)]·3pyridine and of 3a in crystals of
TMA[tBuSalAmi)ZnII]·3.5pyridine (50% probability ellipsoids; hydro-
gen atoms except for amide protons, counterions, and solvent molecules
omitted for clarity).
[((Ad,MeArO)3N)FeII] [δ = 1.04(1) mm/s and ΔEQ = 2.24(1)
mm/s].17 The observed trend in the isomer shift is in good
agreement with the Fe−OAr bond distances in these iron(II)
tris(phenolate) complexes. Only recently, Neese and Petrenko
stated that “the most decisive factor for the isomer shift is the
metal ligand bond”.25 Thus, for the series of complexes, the
shorter the Fe−OAr bond, the lower the isomer shift.
For 1a, the magnetic moment, μeff, of 5.1 μB, obtained by VT
SQUID, is slightly higher than the spin-only value of 4.9 μB but
reproducible and well within the normal range for iron(II)
complexes.
In order to better understand the coordination and electro-
chemistry of iron(II) complex 1a, the corresponding zinc(II)
complex was synthesized by reaction of the deprotonated ligand
with ZnCl2. To ensure that this was a reasonable comparison, the
molecular and crystal structure of TMA[(tBuSalAmi)ZnII]
(TMA-3a; Figure 2) was determined, and it revealed a
coordination geometry very similar to that of 1a. In contrast to
five-coordinate tbp 1a, the zinc complex, TMA-3a, is four-
coordinate trigonal-pyramidal, with an empty axial coordination
site trans to the chelates’ N anchor. The Zn−OAr bond lengths
are 1.916(2) Å with a Zn−Nanchor distance of 2.040(3) Å. The
slight contractions in bond distances are likely due to the smaller
zinc radius and lower coordination number in TMA-3a.
Hydrogen bonding between the amide and phenolate is again
present in this complex.
molecule crystallizes in the Pa3 space group, so the expected
̅
C3 geometry around the metal is crystallographically enforced.
However, the 1H NMR spectra of 1a and 1b suggest that the 3-
fold geometry is maintained in solution as well. The iron center
in 1a is five-coordinate (trigonal-bipyramidal, tbp), with the
tetradentate chelate and an axially bound pyridine. The three
phenolate moieties coordinate the iron center with a small out-
of-plane shift, doop, of 0.026(3) Å [distance of the iron atom
above the plane formed by the three phenolate oxygen (OAr)
atoms] and therefore nearly ideal OAr−Fe−OAr angles of
119.6(2)°. The Fe−OAr distance in 1a is 2.056(3) Å, which is
0.1 Å longer than that in our recently reported iron(II)
tris(phenolate) amine, PNP[((Ad,MeArO)3N)FeII] [d(Fe−
OAr)avg = 1.952(2) Å].17 The anchoring amine is located at a
distance of 2.184(4) Å, again longer than that in our previous
complex, but is still coordinated to the metal center.17 As
expected, the amide protons are oriented toward the metal center
and theoretically available to interact with (and stabilize) a donor
atom of an appropriate terminal ligand. In this complex, the
amide protons are hydrogen-bonded to form a six-membered
ring with the phenolate oxygen atoms. The N(H)−OAr distance
is 2.756(4) Å, and the N−H−OAr angle is 133.0°.
The zero-field 57Fe Moßbauer spectrum of 1a, seen in Figure 2
̈
along with the VT/VF SQUID data, was recorded in solution,
using a 30% 57Fe-enriched sample synthesized from 57FeCl2 in
acetonitrile.24 Both the isomer shift, δ, of 1.19(1) mm/s and the
quadrupole splitting, ΔEQ, of 2.84(1) mm/s are larger than those
for the iron(II) tris(phenolate) amine complex, PNP-
The electrochemistry of Na(18 crown 6)-3a was then studied
and revealed several irreversible, presumably phenolate-based
oxidation events between 0.4 and 1 V, referenced to Fc+/Fc (see
the SI for all electrochemistry). As expected, the redox chemistry
of 1a revealed similar, irreversible ligand-based redox events
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Inorganic Chemistry
Communication
(11) Yeh, C.-Y.; Chang, C. J.; Nocera, D. G. J. Am. Chem. Soc. 2001,
123, 1513−1514.
between 0.6 and 1 V. Two additional redox events were present:
The first event at −1.25 V we assign to an iron(II)/iron(III)
couple that is quasi-reversible at scan rates from 50 to 1600 mV/
s, whereas the origin of the second event at −0.38 V remains as of
yet unknown. As expected, the electrochemistry of the
independently synthesized iron(III) complex, 2a, demonstrated
similar features when investigated.
Because of the quasi-reversible nature of the iron(II)/iron(III)
couple, we believed that the iron(III) species could be
synthesized by chemical oxidation. Chemical oxidation of the
sodium salt of the iron(II) complexes did indeed lead to a
characteristic color change and the isolation of analytically pure
[(RSalAmi)FeIII] (2a, R = tBu, and 2b, R = iPr).
(12) Garner, D. K.; Fitch, S. B.; McAlexander, L. H.; Bezold, L. M.; Arif,
A. M.; Berreau, L. M. J. Am. Chem. Soc. 2002, 124, 9970−9971.
(13) Sahu, S.; Widger, L. R.; Quesne, M. G.; de Visser, S. P.;
Matsumura, H.; Moenne-Loccoz, P.; Siegler, M. A.; Goldberg, D. P. J.
̈
Am. Chem. Soc. 2013, 135, 10590−10593.
(14) Hammes, B. S.; Young, V. G., Jr.; Borovik, A. S. Angew. Chem. Int.
Ed. 1999, 38, 666−669.
(15) Lucas, R. L.; Zart, M. K.; Murkerjee, J.; Sorrell, T. N.; Powell, D.
R.; Borovik, A. S. J. Am. Chem. Soc. 2006, 128, 15476−15489.
(16) MacBeth, C. E.; Gupta, R.; Mitchell-Koch, K. R.; Young, V. G., Jr.;
Lushington, G. H.; Thompson, W. H.; Hendrich, M. P.; Borovik, A. S. J.
Am. Chem. Soc. 2004, 126, 2556−2567.
(17) Kaß, M.; Hohenberger, J.; Adelhardt, M.; Zolnhofer, E. M.;
̈
Further investigation will include a detailed analysis of the
electronic structure of 2a, because initial experiments indicate
that the high-spin FeIII d5 system is unexpectedly complicated
[see the SI for preliminary SQUID, X-band electron para-
Mossin, S.; Heinemann, F. W.; Sutter, J.; Meyer, K., Inorg. Chem.
2013.ASAP. 10.1021/ic4024053
(18) Davis, M. I.; Orville, A. M.; Neese, F.; Zaleski, J. M.; Lipscomb, J.
D.; Solomon, E. I. J. Am. Chem. Soc. 2002, 124, 602−614.
(19) Bandara, D. M. I.; Sono, M.; Bruce, G. S.; Brash, A. R.; Dawson, J.
H. J. Inorg. Biochem. 2011, 105, 1786−1794.
(20) Hwang, J.; Govindaswamy, K.; Koch, S. A. Chem. Commun. 1998,
1667−1668.
(21) Whiteoak, C. J.; Martin, E.; Belmonte, M. M.; Benet-Buchholz, J.;
Kleij, A. W. Adv. Synth. Catal. 2012, 354, 469−476.
(22) Whiteoak, C. J.; Gjoka, B.; Martin, E.; Belmonte, M. M.;
magnetic resonance (EPR), and 57Fe Moßbauer data]. In
̈
addition, iron(II) complexes will be studied for their ability to
reductively activate small molecules, in particular O2.
ASSOCIATED CONTENT
■
S
* Supporting Information
Synthetic details, CHN elemental analyses, 1H/13C NMR, UV−
́
Escudero-Adan, E. C.; Zonta, C.; Licini, G.; Kleij, A. W. Inorg. Chem.
vis, IR, EPR, magnetization, and 57Fe Moßbauer data, as well as
̈
2012, 51, 10639−10649.
X-ray crystallographic details and CIF files. This material is
(23) Licini, G.; Mba, M.; Zonta, C. Dalton Trans. 2009, 5265−5277.
available free of charge via the Internet at http://pubs.acs.org.
(24) Solution-state Moßbauer was necessary; otherwise, difficulty in
̈
completely removing the coordinated solvent from the samples led to
the observation of multiple species.
(25) Neese, F.; Petrenko, T. Quantum Chemistry and Moßbauer
AUTHOR INFORMATION
Corresponding Author
*E-mail: karsten.meyer@fau.de.
Notes
■
̈
Spectroscopy. In Moßbauer Spectroscopy and Transition Metal Chemistry;
̈
Gutlich, P., Bill, E., Trautwein, A., Eds.; Springer: Berlin, 2011; pp 137−
̈
199.
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
This work was supported by the Friedrich-Alexander University
■
Erlangen−Nurnberg (FAU) and DFG. We thank the Bavarian
̈
California Technology Center (BaCaTeC) for generous funding.
M.J.C. thanks the Fulbright Commission for a fellowship. The
authors gratefully acknowledge Dr. Eckhard Bill (MPI CEC,
Mulheim/Ruhr, Germany) for his magnetochemical expertise
̈
and discussions.
REFERENCES
■
(1) Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 6252−6254.
(2) Saouma, C. T.; Peters, J. C. Coord. Chem. Rev. 2011, 255, 920−937.
(3) Scepaniak, J. J.; Fulton, M. D.; Bontchev, R. P.; Duesler, E. N.; Kirk,
M. L.; Smith, J. M. J. Am. Chem. Soc. 2008, 130, 10515−10517.
(4) Vogel, C.; Heinemann, F. W.; Sutter, J.; Anthon, C.; Meyer, K.
Angew. Chem. 2008, 47, 2681−2684.
(5) Scepaniak, J. J.; Vogel, C. S.; Khusniyarov, M. M.; Heinemann, F.
W.; Meyer, K.; Smith, J. M. Science 2011, 331, 1049−1052.
(6) England, J.; Martinho, M.; Farquhar, E. R.; Frisch, J. R.; Bominaar,
E. L.; Munck, E.; Que, L., Jr. Angew. Chem. Int. Ed. 2009, 48, 3622−3626.
(7) Smith, J. M.; Mayberry, D. E.; Margarit, C. G.; Sutter, J.; Wang, H.;
Meyer, K.; Bontchev, R. P. J. Am. Chem. Soc. 2012, 134, 6516−6519.
(8) Bigi, J. P.; Harman, W. H.; Lassalle-Kaiser, B.; Robles, D. M.; Stich,
T. A.; Yano, J.; Britt, R. D.; Chang, C. J. J. Am. Chem. Soc. 2012, 134,
1536−1542.
(9) Berreau, L. M.; Mahapatra, S.; Halfen, J. A.; Young, V. G., Jr.;
Tolman, W. B. Inorg. Chem. 1996, 35, 6339−6342.
(10) Wada, A.; Harata, M.; Hasegawa, K.; Jitsukawa, K.; Masuda, H.;
Mukai, M.; Kitagawa, T.; Einaga, H. Angew. Chem. Int. Ed. 1998, 37,
798−799.
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