Sulfonated diiron complexes as water-soluble models of the [Fe-Fe]-hydrogenase enzyme active site

Article abstract of DOI:10.1021/ic200272x

A series of diiron complexes developed as fundamental models of the two-iron subsite in the [FeFe]-hydrogenase enzyme active site show water-solubility by virtue of a sulfonate group incorporated into the -SCH ₂NRCH₂S- dithiolate uni

Full text of DOI:10.1021/ic200272x

ARTICLE
pubs.acs.org/IC
Sulfonated Diiron Complexes as Water-Soluble Models of
the [FeꢀFe]-Hydrogenase Enzyme Active Site
Michael L. Singleton, Danielle J. Crouthers, Robert P. Duttweiler, III, Joseph H. Reibenspies, and
Marcetta Y. Darensbourg*
Department of Chemistry, Texas A&M University, College Station, Texas 77843, United States
S Supporting Information
b
ABSTRACT: A series of diiron complexes developed as fundamental
models of the two-iron subsite in the [FeFe]-hydrogenase enzyme
active site show water-solubility by virtue of a sulfonate group incorpo-
rated into the ꢀSCH2NRCH2S- dithiolate unit that bridges two
FeI(CO)2L moieties. The sulfanilic acid group imparts even greater
water solubility in the presence of β-cyclodextrin, β-CyD, for which
NMR studies suggest aryl-sulfonate inclusion into the cyclodextrin
cavity as earlier demonstrated in the X-ray crystal structure of 1Na 2 β-
3
CyD clathrate, where 1Na = Naþ(μ-SCH2N(C6H4SO3ꢀ)CH2S-)[Fe-
(CO)3]2, (Singleton et al., J. Am. Chem. Soc. 2010, 132, 8870).
Electrochemical analysis of the complexes for potential as electrocata-
lysts for proton reduction to H2 finds the presence of β-CyD to
diminish response, possibly reflecting inhibition of structural rearrange-
ments required of the diiron unit for a facile catalytic cycle. Advantages of the aryl sulfonate approach include entry into a variety of
water-soluble derivatives from the well-known (μ-SRS)[Fe(CO)3]2 parent biomimetic, that are stable in O2-free aqueous solutions.
’ INTRODUCTION
reduction catalysis capability of the cyanide-containing diiron
complexes, likely because of the sensitivity of these compounds
to both air and acid. There have, however, been several
approaches toward developing more hydrophilic complexes that
exhibit greater stability than the cyanide derivatives.
The synthesis of new small molecule models of the diiron
hydrogenase enzyme active site1 is driven by a desire to better
understand a unique biological system capable of producing
hydrogen at rates comparable to platinum.2 In doing so the
ultimate goal is the development of iron- and sulfur-based
molecular catalysts for use in the cathode of solar powered water
electrolysis cells.3 To date hundreds of these small molecule
models of the form (μ-SCH2XCH2S)[Fe(CO)3ꢀnLn]2 (X =
CR2, NR, O, or S) have been synthesized and their proton
reduction capabilities analyzed through a combination of elec-
trochemical and spectroscopic techniques.4,5 While it might
seem reasonable to conduct these assays in water to match the
ideal environment of the hydrolysis cells, the proton reduction
experiments conducted on the diiron model systems have been
limited almost entirely to organic solvents or mixtures of water
and MeCN. This limitation is due to the inherent insolubility of
the majority of small molecule models in water, which results
from their hydrophobic nature, as well as inherent complexities
with water as solvent in electrochemical studies. In fact, over the
past 10 years only a handful of water-soluble complexes have
been described.
Phosphine ligands have long been used in the [FeꢀFe]-H2ase
model systems as a synthetic substitute for the naturally occurring
CNꢀ ligands.1 The degree of substitution has been shown to play
an important role in fine-tuning the proton reduction capabilities
of hundreds of models of the enzyme active site. Disubstituted
(μ-pdt)[Fe(CO)2L]2 complexes typically reduce protons from
HOAc at the initial FeIFeI/FeIFe0 reduction, but with a large
overpotential.7 Monosubstituted complexes on the other hand
can still reduce protons at the first reduction event but at a
lower overpotential than disubstituted complexes.8 Specialized
ligands such as phosphatriazaadamantane (PTA), tris(hydro-
xymethyl)phosphine, and Ph2P(C6H4SO3ꢀ) have been shown
to increase water solubility in a variety of metal complexes.8ꢀ10 In
the case of PTA, the ligand alone did not impart significant
solubility to (μ-pdt)[Fe(CO)3] [Fe(CO)2TA] in pure water;
however, it did permit studies in mixtures of MeCN/H2O of up to
a 1:3 ratio where the proton reduction capabilities of the com-
plexes were observed to shift to significantly more positive
potentials with increasing amounts of water.8 Sun and co-workers
The dianionic complex (μ-pdt)[Fe(CO)2CN]22ꢀ, (μ-pdt =
-S(CH2)3S-) synthesized from CO substitution in (μ-pdt)[Fe-
(CO)3]2 (1) by CNꢀ, was described as highly water-soluble by
Pickett and co-workers in 1999.6 In the interim there have been
no further reports of work exploring the aqueous, proton
Received: February 8, 2011
Published: April 27, 2011
r
2011 American Chemical Society
5015
dx.doi.org/10.1021/ic200272x Inorg. Chem. 2011, 50, 5015–5026
|

Inorganic Chemistry
ARTICLE
Figure 1. (A) Cartoon of X-ray crystal structure of sulfonated diiron hexacarbonyl complex, 1A, encapsulated within two β-cyclodextrins.12 (B) Diagram
of R-, β-, and γ-cyclodextrins.13 (C) Structure of ferrocene in β-cyclodextrin.14 (D) Cartoon of the O2-binding iron-porphyrin biomimetic cyclodextrin
ꢀ 16
system developed by Kano and co-workers. R = -p-C6H4SO3
.
further used the di-N-acetylated phosphatriazaadamantane
(DAPTA) ligand to impart moderate solubility of ((μ-pdt)[Fe-
(CO)2L]2 in water which allowed the first electrochemical studies
in 100% water (0.05 M n-Bu4NPF6 solution).9 These studies
indicated that water played a significant role in altering both the
reduction potential and the electrocatalytic response of the model
complexes, with the reduction of protons occurring more positive
than the initial FeIFeI/FeIFe0 reduction as a result of the
protonation of the amines in the PTA and DAPTA ligands.
While successful for developing less hydrophobic diiron model
systems, the ligand-based solubilization technique limits the variety
of complexes that might be synthesized. In contrast, incorporation
of a water-solubilizing group into the dithiolate linker appears to
offer more versatility. By modifying D-glucopyranoside to include a
dithiolate unit Weigand and co-workers were able to obtain sugar
functionalized diiron models.11 These species, initially synthesized
as the protected tetra-O-acetyl derivatives, showed greatly in-
creased hydrophilicity upon deprotection of the hydroxyl units.
Unfortunately, the deprotected species was not stable in protic
solvents, thus limiting its applicability.
Recently, we described the synthesis of a diiron model complex
containing an aryl sulfonate modified azadithiolate linker, Naþ
(μ-SCH2N(C6H4SO3ꢀ)CH2S-)[Fe(CO)3]2, 1Na.12 While the
addition of this moiety into the parent hexacarbonyl complex
renders it insoluble in most nonpolar organic solvents, it imparts
moderate solubility (∼1.1 mM) in pure water. Additionally, the
complex remains stable for an extended period of time in
deaereated water (up to several months when stored in the
refrigerator), allowing study of the diiron complex in water by
NMR, IR, UVꢀvis, and electrochemical methods. Addition-
ally, the aryl sulfonate group promoted inclusion of the
complex in β-cyclodextin leading to the structural character-
the cyclodextrin greatly increased the water solubility of 1Na
and opened the possibility for using the inclusion in the cavity
to mimic the protein environment of the enzyme.
Cyclodextrins, cyclic oligosacharides consisting most com-
monly of 6ꢀ8 glucose units, have properties long recognized as
ideal for biomimetics, Figure 1B.13 Their hydrophobic cavities,
ranging in diameter from 4.5 to 9.5 Å, depending on the number
of glucose units, and having a height of ∼8.0 Å, have in fact been
demonstrated to serve as hosts for organometallics.14 The
hydrophilic hydroxyl rims provide hydrogen bonding sites.15
An important example is the inclusion of ferrocene in R-, β-, and
γ- cyclodextrins, the adducts of which have all been structurally
characterized, Figure 1C . In all three adducts, the cyclodextrins
incorporated at least one of the cyclopentadienyl rings of
ferrocene, affecting the spectroscopic and electrochemical prop-
erties of the organometallic unit.
Furthermore inclusion in the cyclodextrins can provide similar
site isolation as active site cavities in proteins. A recent example of
this application, the one that was inspirational for our work, was
reported by Kano and co-workers, Figure 1D.16 Through the
inclusion of an iron tetrakis(4-sulfonatophenyl)porphyrin inside
a dimeric cyclodextrin system, where the two cyclodextrins were
connected at the 2 position via a pyridine linker, they showed that
the cyclodextrin super structure provided a high degree of site
isolation and allowed the iron porphyrin to bind oxygen in water
without forming the commonly observed μ-oxo dimers. While
the intriguing cyclodextrin-containing biomimetic models such
as this are not common, the inherent inclusion properties of
cyclodextrins and the extensive literature on methods to modify
them make them attractive targets for exploring the effects of
inclusion in a cavity on models of [FeꢀFe]-hydrogenase.
In the present work we have used the sulfonated hexacarbonyl
complex, 1Na, as a precursor to generate a series of phosphine-
substituted complexes of the type Na(μ-SCH2N(C6H4SO3)-
CH2S-)[Fe(CO)3][Fe(CO)2L] (L = P(OMe)3 (2Na); PTA
ization of a 1Na 2 β-cyclodextrin clathrate, which showed the
3
aryl sulfonate group penetrating through the cavity of one of
the cyclodextrins, Figure 1A. As expected, the interaction with
5016
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
(3Na); and PPh3 (4Na) and Na(μ-SCH2N(C6H4SO3)CH2-
S-)[Fe(CO)2PMe3]2 (5Na), that have varying degrees of water
solubility. Through structural analysis and electrochemical stud-
ies in MeCN we show that the sulfonate group can impart
modest water solubility to complexes 2Na, 3Na, and 5Na
without drastically altering their structural or electrochemical
properties. The addition of the large hydrophobic PPh3 ligand,
however, overwhelms the effects of the sulfonate group and
renders the model complex insoluble in H2O.
electrode was used and the supporting electrolyte solution was 0.1 M
Et4NBF4. All samples in MeCN were run at a concentration of 2 mM,
and potentials are reported relative to the Fc/Fcþ couple as 0.00 V.
Water solubility studies of the diiron complexes were performed
using a Shimadzu UV-2450 UVꢀvis spectrometer. Deaerated water was
added to known amounts of each complex until no solid material was
visible in solution. This stock solution was then used to generate several
samples of known concentrations allowing the molar absorptivities for
each absorption band in the UVꢀvis spectra to be calculated. Maximum
concentrations were obtained by adding enough of each complex to
20 mL of deaerated water that solid remained even after sonicating each
sample for 10 min. The slurry was then centrifuged for 30 min or until no
suspended solid was visible in solution. The UVꢀvis spectrum of the
supernatant was recorded, and Beer’s law was used to calculate the
concentration.
X-ray Diffraction Studies. For complexes 2TEA and 5TEA low-
temperature (110 K) X-ray data were obtained on a Bruker Apex-II CCD
based diffractometer (Texas A&M University) (Mo sealed X-ray tube,
KR = 0.71073 Å). For complexes 1Na, 3TEA, and 4 TEA low-temperature
(110 K) X-ray data were obtained on a Bruker-D8 Adv GADDS general-
purpose three-circle X-ray diffractometer (Texas A&M University) (Cu
sealed X-ray tube, KR = 1.54184 Å). Space groups were determined on the
basis of systematic absences and intensity statistics, and structures were
solved by direct methods and refined by full-matrix least-squares on F2. All
non-hydrogen atoms were refined with anisotropic thermal parameters. H
atoms were placed at idealized positions and refined with fixed isotropic
displacement parameters; anisotropic displacement parameters were
employed for all non-hydrogen atoms. Complexes 1Na and 2TEA
were found as racemic twins (BASF = 0.23126 and 0.45113,
respectively). Solvent molecules in 2TEA and 5TEA presented with
significant disordering and were SQUEEZEd out using Platon.18 The
following programs were used: data collection, APEX219a and GADDS,19b
data reduction, SAINT19c; absorption correction and twin refine-
ment, SADABS19d and TWINABS;19d cell refinement SHELXTL;19d
structure solutions, SHELXS-9719d (Sheldrick); and structure refine-
ment, SHELXL-9719d (Sheldrick). The final data presentation and
structure plots were generated in X-Seed Version 1.5.20
Additionally, to expand on our previous study,12 the inter-
action of the complexes with cyclodextrins is studied in H2O
through NMR spectroscopy and cyclic voltammetry, to deter-
mine the effect of cyclodextrin on the phosphine-substituted
complexes. All of the complexes show distinct changes in their
proton resonances, with 2Na and 3Na behaving similarly to 1Na.
The 5Na complex on the other hand shows different changes,
and the specific interaction with cyclodextrin is not as well-
defined. For complex 4Na, the aqueous solubility is only slightly
improved in the presence of cyclodextrin, limiting the study of
the interaction by NMR.
For the electrochemical studies, our initial hypothesis was that,
because of the generally weaker FeꢀFe bond in substituted
complexes and the possibility of hydrogen bonding interactions
in the P(OMe)3 and PTA derivatives, the cyclodextrins might
induce changes in the electrochemical properties of the model
complexes by influencing the variety of conformations proposed
to occur in the electrocatalytic cycle.17 As will be shown, it
appears that the effect of the cyclodextrin is to hinder proton
reduction rather than facilitate it, likely by (a) preventing the
structural rearrangements needed in the catalytic cycle; (b)
by inhibiting proton delivery processes within the cavity; or
(c) both.
’ EXPERIMENTAL SECTION
General Experimental Procedures. All solvents used were
purified and degassed via a Bruker solvent system. Unless otherwise
stated, all syntheses and manipulations were performed using standard
Schlenk-line and syringe/rubber septa techniques under N2 or in an
argon atmosphere glovebox. Complexes 1Na and 1TEA were synthe-
sized according to recently published procedures.12 β-cyclodextrin was
purchased from TCI America and used without further purification. All
other reagents were purchased from Aldrich Chemical Co. and used as
received. High resolution electrospray ionization mass spectrometry
(HRMS) analysis was performed by the Laboratory for Biological Mass
Spectrometry at Texas A&M University. Infrared spectra were recorded
on a Bruker Tensor 37 spectrometer in either CaF2 or NaCl solution
cells with a 0.1 mm path length. 1H NMR spectra were acquired on an
Inova 500 spectrometer operating at 500 MHz. All NMR samples were
’ SYNTHESES
Crystallization of Na(μ-SCH2N(C6H4SO3)CH2S-)[Fe(CO)3]2
(1Na). 1Na was synthesized according to the recently described
procedure.12 Wet 1Na was dissolved to saturation in THF.
Crystals of X-ray quality were obtained by cooling this solution
at 5 °C overnight.
Synthesis of Na(μ-SCH2N(C6H4SO3)CH2S-)[Fe(CO)3][Fe-
(CO)2P(OMe)3], (2Na). To a 100 mL Schlenk flask containing
1Na (0.250 g, 0.443 mmol) and trimethylamine-N-oxide
(0.050 g, 0.066 mmol) was added MeCN (20 mL) resulting in
a dark red solution. After stirring for 10 m at 22 °C, P(OMe)3
(0.10 mL 0.090 mmol) was added. IR analysis of the solution
showed bands as follows: ν(CO) in MeCN: 2050 m, 1994 s,
1973 sh, 1940 w. The reaction was stirred for 30 m after which
time the solvent was removed in vacuo. The resulting solid was
completely dissolved in tetrahydrofuran (THF) and loaded onto
the top of a silica column (4 cm ꢁ20 cm). The dark red product
was eluted with MeOH, and solvent was removed with a rotary
evaporator to yield 0.115 g (0.174 mmol, 39.2%) of the pure
product. The Et4Nþ salt (2TEA) was prepared in an analogous
manner starting from 1TEA. Crystals of X-ray quality were
obtained by slow diffusion of ether into a CH2Cl2 solution
of 2TEA. IR (THF) ν(CO): 2048 (m); 1992 (s); 1971 (m);
1942 (w), cmꢀ1; 1H NMR (DMSO): 3.72 (d, JHP = 11.2, 9H);
1
prepared in either D2O or DMSO. H chemical shifts are reported in
parts per million relative to H2O (4.79 ppm) or DMSO (2.50 ppm).
Voltammograms were obtained using a standard three electrode cell
under argon at room temperature using a Bioanalytical System (BAS)
100 electrochemical workstation with a glassy carbon working electrode
and a platinum wire auxiliary electrode. For aqueous work, a saturated
KCl Ag/AgCl reference electrode was used, and the supporting electro-
lyte solution was 0.01 M NaCl. Samples were run at a concentration of
∼0.1 mM because of their low solubility in the NaCl buffer solution.
Between each scan the glassy carbon electrode was removed and
polished using a 1 μm polycrystalline diamond suspension and rinsed
with both acetone and deionized water to remove any adsorbed material.
All aqueous potentials are reported relative to the Ag/AgCl couple as
0.00 V. For work done in MeCN, a saturated Ag/AgNO3 reference
5017
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Scheme 1. Synthesis of Phosphine Substituted Complexes 2Naꢀ5Na: (A) (1) Me3NO, MeCN (2) P(OMe)3; (B) PTA, MeOH,
Δ; (C) (1) Me3NO, MeCN (2) PPh3; (D) PMe3, THF, Δ
4.35 (d, J = 10.8 Hz, 2H); 4.38 (d, J = 10.8 Hz, 2H) ppm; 6.79 (d,
J = 8.3 Hz, 2H); 7.46 (d, J = 8.3, 2H); Electronic absorption
spectrum (H2O): λmax (ε, Mꢀ1 cmꢀ1) 483 (1450), 335 (8230),
308 (11800), 274 (13310), HRMS (ESI) m/z calcd for
C16H17Fe2NO11PS3 Theoretical Mass: 637.8400 g/mol;
Found Mass: 637.8371 g/mol.
IR analysis of the solution showed bands as follows: ν(CO) in
MeCN: 2048 m, 1990 sh, 1931 w. The reaction was allowed to
stir for 30 min, after which time the solvent was removed in
vacuo. The resulting red solid was completely dissolved in THF
and loaded onto the top of a silica column (4 cm ꢁ 20 cm). The
dark red product was eluted with MeOH, and solvent was
removed with a rotary evaporator. The dried residue was
triturated with H2O and dried under vacuum to yield 0.049 g
(0.061 mmol, 34.5%) of the pure product. The Et4Nþ salt
(4TEA) was obtained in an analogous manner starting from
1TEA. Crystals of X-ray quality were obtained by cooling a
saturated MeOH solution of 4TEA at ꢀ5 °C for 2 weeks. IR
1H NMR (DMSO): 4.28 (d, J = 12.7, 2H); 6.65 (d, J = 8.8 Hz,
2H); 7.40 (d, J = 8.8 Hz, 2H); 7.47 (m, 3H); 7.55 (m, 6H); 7.63
(m, 6H) ppm; Electronic absorption spectrum (H2O): λmax
(ε, Mꢀ1 cmꢀ1) 489 (1260), 330 (9330), 303 (14210), 229
(15130), HRMS (ESI) m/z calcd for C31H23Fe2NO8PS31- Theo-
retical Mass: 775.9022 g/mol; Found Mass: 775.8988 g/mol.
Synthesis of Na(μ-SCH2N(C6H4SO3)CH2S-)[Fe(CO)2PMe3],
(5Na). To a slurry of 1Na,(0.250 g, 0.44 mmol) in THF (20 mL)
was added PMe3, (0.5 mL, 4.2 mmol). The solution was then
heated to 60 °C, and the solution was monitored by IR
spectroscopy. After 4ꢀ6 h no ν(CO) bands were visible above
2000 cmꢀ1, and the reaction was stopped. The resulting dark red
solution was concentrated to ∼5ꢀ10 mL, and ether (50 mL) was
added resulting in a red-purple precipitate. The solid was filtered
under N2 and rinsed on the filter with ether (30 mL), toluene
(30 mL), and hexanes (30 mL). After drying under vacuum for
24 h, 0.181 g (0.273 mmol, 61.6%) of 5Na was obtained. The
Et4Nþ salt (5TEA) was prepared in an analogous manner
starting from 1TEA. Crystals of X-ray quality were obtained by
slow diffusion of ether into a CH2Cl2 solution of 5TEA. IR
(THF) ν(CO): 1982 (w); 1948 (s); 1903 (m), cmꢀ1; 1H NMR
(DMSO): 1.47 (d, JHP = 9.28, 18H); 4.17 (s, 4H); 6.69 (d, J =
8.3 Hz, 2H); 7.39 (d, J = 8.3 Hz, 2H) ppm; Electronic absorption
spectrum (H2O): λmax (ε, Mꢀ1 cmꢀ1) 517 (370), 352 (1300),
1-
Synthesis of Na(μ-SCH2N(C6H4SO3)CH2S-)[Fe(CO)3][Fe(CO)2-
PTA], (3Na). To a slurry of 1Na, (0.200 g 0.355 mmol) in MeOH
(20 mL) was added PTA, (0.111 g, 0.704 mmol). The solution was
then heated to 60 °C, and the reaction was monitored by IR
spectroscopy. The reaction mixture was stirred overnight (∼20 h)
after which time IR analysis of the solution indicated that it contained
primarily the monosubstituted diiron complex as well as a small
amount of the starting material. The solvent was removed in vacuo,
and the residue was dissolved in THF and loaded onto the top of a
silica column (4 cm ꢁ20 cm). Elution with methanol gave two bands
which contained, in order, the unreacted starting material and then
the product 3Na. The fractions containing the second band were
concentrated to ∼3ꢀ5 mL on a rotary evaporator and ether was
added resulting in a red precipitate. The solid was dried under
vacuum to yield 0.108 g (0.156 mmol, 43.8%) of 3Na. The Et4Nþ
salt (3TEA) was synthesized by counterion exchange using Et4NCl
in CH2Cl2. Crystals of X-ray quality were obtained by slow diffusion
of ether into a concentrated CH2Cl2 solution of 3TEA. IR (THF)
ν(CO): 2042 (m); 1986 (s); 1967 (sh); 1930 (w) cmꢀ1; 1H NMR
(DMSO): 4.05 (s, 6H); 4.21 (d, J = 13.2 Hz, 2H); 4.38 (d, J = 12.7
Hz, 3H); 4.46 (d, J = 12.7 Hz, 3H); 4.51 (d, J = 13.2 Hz, 2H); 6.76
(d, J = 8.8 Hz, 2H); 7.44 (d, J = 8.8 Hz, 2H) ppm; Electronic
absorption spectrum (H2O): λmax (ε, Mꢀ1 cmꢀ1) 500 (540), 346
(2880), 275 (5560); HRMS (ESI) m/z calcd for C19H20Fe2-
(THF) ν(CO): 2046 (m); 1990 (s); 1975 (m); 1936 (w), cmꢀ1
;
1-
N4O8PS3 Theoretical Mass: 670.8880 g/mol; Found Mass:
670.8847 g/mol.
Synthesis of Na(μ-SCH2N(C6H4SO3)CH2S-)[Fe(CO)3][Fe(CO)2-
PPh3], (4Na). To a 100 mL Schlenk flask containing 1Na
(0.100 g, 0.177 mmol), trimethylamine-N-oxide (0.015 g,
0.200 mmol), and PPh3 (0.051 g, 0.195 mmol) was added
MeCN (20 mL), and the solution was heated to 60 °C.
5018
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
274 (3040), 247 (3400), HRMS (ESI) m/z calcd for C18H26Fe2-
NO8P2S3 Theoretical Mass: 637.9045 g/mol; Found Mass:
637.9071 g/mol.
by heating a solution of 1Na and multiple equivalents of the PTA
ligand overnight. All of the monosubstituted complexes are ob-
tained as air-stable red solids in moderate yield and can be isolated
as the Naþ salt as depicted in Scheme 1 or as the Et4Nþ salt by
direct synthesis from 1TEA or by counterion exchange using
1-
’ RESULTS AND DISCUSSION
1
Et4NCl. H NMR analysis of the complexes shows two signals
Synthesis of Na(μ-SCH2N(p-C6H4SO3ꢀ)CH2S)[Fe(CO)3][Fe-
(CO)2L], L = P(OMe)3 (2Na); PTA (3Na); or PPh3 (4Na) and
Na(μ-SCH2N(p-C6H4SO3ꢀ)CH2S)[Fe(CO)2PMe3]2 (5Na). The
condensation reaction between formaldehyde, an amine, and
(μ-SH)2[Fe(CO)3]2 to give (μ-SCH2N(R)CH2S)[Fe(CO)3]2 as
reported by Rauchfuss and co-workers has become a widely used
procedure for the incorporation of useful R groups into the dithiolate
linker of diiron complexes as parent models of the [FeꢀFe]-
Hydrogenase enzyme active site.21 We have recently demonstrated
that sulfanilic acid can be used in this method to introduce an aryl
sulfonate group into the dithiolate linker of the model complexes to
give the anionic (μ-SCH2N(p-C6H4SO3ꢀ)CH2S)[Fe(CO)3]2.12
This compound acts as precursor to a series of water-soluble
phosphine-substituted complexes through the synthetic routes
shown in Scheme 1.
for the methylene protons in the dithiolate linker of the monopho-
sphine complexes, resulting from the loss of symmetry between the
two irons upon substitution of a single CO. Table 1 lists the IR
stretching frequencies in THF for the complexes shown in
Scheme 1. As expected, ν(CO) stretching frequencies for the
monosubstituted complexes are lower than those of 1Na, consis-
tent with the increased donation into the Fe atoms by the
phosphine ligands versus CO; the increased electron density on
Fe based on these ν(CO) values is in the order 3Na > 4Na > 2Na.
The reaction of 1Na with excess PMe3 at 60 °C in THF gives
the disubstituted complex 5Na in ∼60% yield. IR analysis of the
product shows significantly lower ν(CO) stretching frequencies
for 5Na as compared to the monosubstituted complexes, con-
sistent with the higher degree of substitution. 1H NMR data find
a single signal for the four methylene protons in the SRS linker at
room temperature indicating that one phosphine has substituted
onto each iron. Additionally, this single signal indicates that the
sulfanilate moiety is flipping rapidly between both sides as is seen
in the case of the (μ-pdt) analogue.22 Unlike the stable mono-
substituted complexes, 5Na is air sensitive and degrades rapidly
in solutions exposed to air. As a solid, 5Na is air tolerant for short
amounts of time; however, prolonged exposure also leads to
decomposition.
Treatment of 1Na with 1 equiv of Me3NO followed by either
P(OMe)3 or PPh3 yielded the monosubstituted complexes 2Na
and 4Na, respectively. The PTA derivative, 3Na, was synthesized
Table 1. IR Stretching Frequencies in the ν(CO) Region in
THF for Complexes 1Naꢀ5Na
complex
ν(CO) cmꢀ1
1Naa
2Na
3Na
4Na.
5Na
2074 (w); 2035 (s); 1996 (m)
Water Solubility of 2Naꢀ5Na. The sulfonate group endows a
degree of water solubility on (μ-SCH2N(C6H4SO3ꢀ)CH2-
S-)[Fe(CO)2L]2 (L = CO or PR3) that varies depending on
the ligands. In the case of the parent hexacarbonyl complex, 1Na,
the maximum obtainable concentration was ∼1.1 mM at 22 °C.
Upon substitution, the solubilities vary depending on the nature
2048 (m); 1992 (s); 1971 (m); 1942 (w)
2042 (m); 1986 (s); 1967 (sh); 1930 (w)
2046 (m); 1990 (s); 1975 (m); 1936 (w)
1982 (w); 1948 (s); 1903 (m)
a From ref 12; (s) strong; (m) medium; (w) weak; (sh) shoulder.
Table 2. Crystallographic Experimental Data for the Complexes
1Na THF
2TEA
3TEA
4TEA
5TEA
3
formula
formula weight
C18H18Fe2NNaO11S3
655.20
C24H37Fe2N2O11PS3
768.41
C27H40Fe2N5O8PS3
801.49
C39H43Fe2N2O8PS3
906.60
C27H46Fe2N2O8P2S3
796.48
temperature (K)
wavelength (Å)
Z
110(2)
1.54178
2
110(2)
0.71073
4
110(2)
1.54178
4
110(2)
1.54178
2
110(2)
0.71073
4
Dcalc’d (Mg/cm3)
1.753
12.436
triclinic
P1
0.671
1.527
1.455
7.838
triclinic
P1
1.440
μ (mmꢀ1
)
0.509
9.241
1.092
crystal system
space group
unit cell
a (Å)
orthorhombic
Pca21
monoclinic
P2/n
monoclinic
P21/c
6.4151(9)
12.6371(9)
15.6834(16)
96.703(2)
95.498(9)
98.239(10)
1241.5(2)
1.045
14.046(14)
33.262(15)
16.28(3)
90.00
14.968(2)
7.6777(11)
30.718(4)
90.00
10.1101(18)
11.2927(19)
18.607(3)
10.311(2)
34.159(7)
10.437(2)
90
b (Å)
c (Å)
R
102.336(6)
93.946(6)
β
90.00
99.042(7)
90.00
91.713(2)
90
γ
90.00
90.895(6)
volume (Å3)
7606(18)
0.962
3486.3(9)
1.030
2069.4(6)
3674.4(12)
1.065
goodness-of-fit
1.029
a
b
R1 , wR2 (%) [I > 2σ(I)]
0.0840, 0.2325
0.0738, 0.1362
0.0806, 0.1543
0.1598, 0.1828
0.0521, 0.0982
0.0937, 0.1080
0.0428, 0.1036
0.0560, 0.1094
a
b
R1 , wR2 (%) all data
0.1020, 0.2550
2
0.1649, 0.1692
2 2 1/2
a R1 = ∑||Fo| ꢀ |Fc||/∑|Fo|. b wR2 = [∑w(Fo ꢀ Fc2)2/∑w(Fo ) ]
.
5019
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Figure 2. (Left) Extended structure of 1Na as a stick model showing the sodium ion-controlled array in the solid state. Protons have been omitted for
clarity and sodium ions are represented as balls. (Right) Close up view of the Naþ ions showing the coordination environment around them. All protons
except the water protons have been removed for clarity.
Figure 3. Solid states structures of (A) 2TEA, (B) 3TEA, (C) 4TEA, and (D) 5TEA as ball and stick models. Protons and the Et4Nþ counterions have
been omitted for clarity.
5020
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Table 3. Comparison of Selected Metric Data for 1Na,
1TEA,12 and 123
Table 4. Selected Metric Data for 2TEAꢀ5TEA
2TEA
3TEA
4TEA
5TEA
123
1TEA12
1Na
FeꢀFe
2.504(3) 2.532(2) 2.506(1) 2.600(1)
FeꢀFe
2.510(1)
1.800
2.498(1)
1.797
2.512(6)
1.786
FeꢀCavg
1.836
1.106
1.759
1.169
1.768
1.143
1.763
1.154
FeꢀCavg
CꢀOavg.
CꢀOavg.
1.136
1.140
1.155
Fe(1)ꢀS(1)
Fe(1)ꢀS(2)
Fe(2)ꢀS(1)
Fe(2)ꢀS(2)
S(1)ꢀC(7)
S(2)ꢀC(8)
2.572(3) 2.260(2) 2.274(1) 2.250(1)
2.571(2) 2.281(3) 2.269(1) 2.269(1)
2.186(3) 2.255(3) 2.271(1) 2.252(1)
2.170(2) 2.256(2) 2.281(1) 2.257(1)
1.984(7) 1.985(8) 1.852(5) 1.857(3)
1.984(8) 1.982(8) 1.854(5) 1.852(3)
Fe(1)ꢀS(1)
Fe(1)ꢀS(2)
Fe(2)ꢀS(1)
Fe(2)ꢀS(2)
S(1)ꢀC(7)
S(2)ꢀC(8)
2.254(1)
2.249(1)
2.254(1)
2.249(1)
1.823(4)
1.818(4)
2.251(1)
2.269(1)
2.257(1)
2.278(1)
1.860(5)
1.861(5)
2.259(8)
2.282(7)
2.280(8)
2.275(8)
1.86(3)
1.79(3)
Fe(1)ꢀS(1)ꢀFe(2)
Fe(1)ꢀS(2)ꢀFe(2)
Fe(2)ꢀFe(1)ꢀC(1)
Fe(1)ꢀFe(2)ꢀLapical
67.3(1)
68.3(1)
68.23(8) 66.92(4) 70.57(3)
67.86(7) 66.83(4) 70.35(3)
Fe(1)ꢀS(1)ꢀFe(2)
Fe(1)ꢀS(2)ꢀFe(2)
Fe(2)ꢀFe(1)ꢀC(1)
Fe(1)ꢀFe(2)ꢀC(6)
67.67(4)
67.85(4)
148.3
66.97(4)
67.02(4)
147.4(1)
155.8(2)
66.8(2)
67.3(2)
155.7(9)
148.1(1)
154.1(3) 152.6(3) 156.1(2) 152.6(1)
149.0(1) 148.9(3) 156.19(5) 142.7(1)
148.3
C(1)ꢀFe(1)ꢀFe(2)ꢀLapical 5.0(7)
10.5(8)
9.8(5)
11.3(3)
C(1)ꢀFe(1)ꢀFe(2)ꢀC(6)
0
7.9(5)
25.7(1)
angle between the apical carbonyl units for 1TEA. However,
when the Et4Nþ counterion is exchanged for Naþ, this torsion
angle increases by ∼18° (25° > 1). In the [FeꢀFe]-H2ase
enzyme active site, a number of dipoleꢀdipole and hydrogen
bonding interactions are purported to stabilize the unique
geometry and semibridging carbonyl.1 The possibility that an
ion dipole interaction in the model complexes could be used to
enforce structural changes is thus intriguing.24 Nevertheless, UV
spectroscopic studies showed no discernible difference between
1TEA and 1Na in the solvents used in this study indicating that
these interactions in 1Na exist only in the confines of the
solid state.
Despite repeated attempts, crystals of X-ray quality were not
obtained for the Naþ salts of the substituted complexes; however,
the respective Et4Nþ salts (2TEAꢀ5TEA) of the compounds
readily crystallize from either slow diffusion of ether into a CH2Cl2
solution of the compounds or by cooling saturated MeOH
solutions. The core structure of all four molecules contains the
familiar 2Fe2S butterfly motif common to this class of com-
pounds.25 Additionally, both irons have a square based pyramidal
geometry with the two sulfurs defining the shared edge between
them. As is the normal case with the structures of most FeIFeI
model complexes the CO ligands are all terminal. The pyramida-
lization of the N in the dithiolate linker is very modest (ca. 0.1 to 0.2
Å above the C3 plane around it) throughout the series of complexes
and is similar to that observed for 1Na and the corresponding
aniline derivative which lacks the sulfonate group.21
The solid state structures of the monosubstituted complexes
are consistent with the degrees of substitution indicated by the
spectral data and are overall structurally similar to the analogous
propanedithiolate complexes.8,26 For 2TEAꢀ4TEA the phos-
phine ligands are situated on the iron distal to the aryl sulfonate
group, likely to minimize steric repulsions between the ligand
and the aryl group of the dithiolate. While the weaker donor
ligands, P(OMe)3 and PPh3, adopt an apical position within their
respective FeS2(CO)2L square pyramids, the PTA ligand as-
sumes a basal orientation consistent with its stronger donor
ability and small size. The FeꢀFe bond distances for 2TEAꢀ
4TEA remain similar throughout the series, Table 4, with only an
∼0.03 Å change from the weaker donor-containing complexes
of the ligands. Both complexes 2Na and 5Na show slightly higher
maximum concentrations, about 1.4ꢀ1.9 mM at 22 °C, than
complex 1Na. For complex 4Na, the hydrophobic PPh3 ligand
dramatically reduces solubility and lowers the maximum con-
centration to <0.1 mM, making it the least water-soluble of the
complexes. Conversely when the more hydrophilic PTA ligand is
used the maximum concentration is greatly increased, with 3Na
dissolution of up to ∼3 mM in pure water. The significantly
greater water solubility of the 3Na complex makes it one of the
most water-soluble diiron complexes reported to date. The
CNꢀ-containing complexes are expected to be even more
water-soluble based on the original account by Pickett and co-
workers;6 however, no quantitative data on their solubilities was
reported. It should be noted that the solubilities for all of the
complexes were measured in pure water; the addition of small
amounts of salt (NaCl, KCl, or Et4NCl) results in partial
precipitation of the complexes from solution.
Table 2 lists the crystal data for the five structures. While the
structure of 1TEA was recently reported, that of the sodium salt,
1Na, was not.12 Upon cooling a THF solution of 1Na at 5 °C
overnight in the presence of a small amount of water, large X-ray
quality crystals grew. While similar to the Et4Nþ structure, there
are key differences between the two salts. The extended packing
diagram of 1Na, Figure 3, shows that two Naþ ions interact with
adjacent sulfonate groups and bridge between pairs of the diiron
moieties. In addition to a THF solvent molecule that also binds
to the Naþ, two water molecules connect each Naþ to the Naþ of
an adjacent dimer to create a two-dimensional array of the diiron
molecules in the solid state. Though not within bonding distance
to the Naþ, an apical carbonyl O atom from an adjacent array at
2.933 Å blocks the sixth coordination site on the Naþ leaving it
with a pseudosquare pyramidal geometry. A long-range ion-
dipole interaction between the Naþ and the carbonyl oxygen
could be the explanation for the increased C1ꢀFe1ꢀFe2-C6
torsion angle of 1Na.
X-ray Diffraction Studies of 1Na and 2TEAꢀ5TEA. Table 3
lists selected metric parameters of (μ-pdt)[Fe(CO)3]2 (1),
1TEA, and 1Na. In 123 the Fe(CO)3 units are eclipsed, and
there is only a slight increase in the C1ꢀFe1ꢀFe2ꢀC6 torsion
5021
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Figure 4. Cyclic voltammograms of complexes 2Na, 3Na, 4Na, and 5Na in MeCN with 0.1 M Bu4NBF4 at a scan rate of 200 mV/s. Concentration of all
samples is 2 mM.
2TEA and 4TEA to complex 3TEA; they fall within the typical
range for FeꢀFe bond distances found in the diiron model
complexes, 2.49ꢀ2.56.27 The similarities to previously reported
systems means that the sulfonate group, while altering the
solubilities, is far enough removed from the diiron unit to have
influence on significant changes in the core structure and reac-
tivity of the complexes.
The structure of 5TEA shows similarities to both the monosub-
stituted complexes and its (μ-pdt) analogue. Both of the PMe3
ligands adopt basal positions in the FeS2(CO)2L square pyramids
and are transoid to each other in the diiron structure. The core
geometry and twisting between the two iron units, defined by the
C1ꢀFe1ꢀFe2ꢀC6 torsion angle, is only slightly larger than in the
monosubstituted species. However, the FeꢀFe bond length of
2.600(1) Å in 5TEA is significantly longer than most model
complexes and is instead closer to the values reported for the enzyme
active site (∼2.6).28 This value is also one of the largest reported for a
FeIFeI species, even among other disubstituted complexes.27
Interaction with β-Cyclodextrin. In addition to providing
water solubility, the aryl sulfonate group also promotes interaction
between the model complexes and β-cyclodextrin (β-CyD). Our
minor band shifts are observed: 2079, 2042 (s), 2008 (s), see
Supporting Information, Figure S8.
Nuclear Magnetic Resonance (NMR) studies on the interac-
tion of cyclodextrins with guest molecules find greater changes
for the inward directed C3 and C5 protons on the cyclodextrin as
compared to the other protons.29 While the direction and
magnitude of this change is dependent on the nature of the
group inside the cyclodextrin cavity, aromatic groups typically
result in an increased shielding of the C3 and C5 protons and
thus an upfield shift. An upfield shift is also generally observed for
the protons on the aromatic guest molecule. This was the case
observed for 1Na and these same resonances were used in
studying the substituted derivatives. The NMR studies, con-
ducted in D2O and focused on the dependence of both the
1
phosphine substituted complexes and the cyclodextrin H che-
mical shifts on the ratio of the two components, indicate that all
of the substituted complexes can interact with cyclodextrin (see
Supporting Information); however, the specific nature of this
interaction varies throughout the series.
For complex 2Na, as the ratio of β-CyD to 2Na is increased,
there is an upfield shift of the aromatic proton resonances of the
aryl sulfonate group from 7.68 and 6.92 ppm for pure 2Na to 7.51
and 6.79 ppm at a CyD/2Na ratio of 9:1. Additionally, moderate
shifts in the C3 and C5 proton resonances of the cyclodextrin are
observed when the ratio is reversed to 1:9 of CyD/2Na. Shifts of
this magnitude are consistent with an interaction between the
iron complex and the cyclodextrin and are similar to those
observed for 1Na.12 Additionally because the changes in chemi-
cal shift for the cyclodextrin occur primarily for the C3 and C5
protons, which are oriented toward the interior cavity of the
cyclodextrin, the behavior is consistent with inclusion in the
cavity of the cyclodextrin. Similar effects are observed for 3Na as
recent description of the 1Na 2β-CyD solid state structure,
3
Figure 1A, showed that the aryl sulfonate group of the complex
penetrates the cavity, putting the aromatic ring inside of the
cyclodextrin with the sulfonate group protruding through the
primary hydroxyl rim of the cyclodextrin.12 Interestingly, the crystal
structure also showed that a second cyclodextrin could interact with
the Fe(CO)3 unit opposite the aryl sulfonate, though no evidence
of such an interaction was observed in solution. The ν(CO) IR
spectra of the all-CO parent sulfonate, 1 Na in water shows a band
pattern and positions typical of all (μ-SRS)[Fe(CO)3]2 complexes,
2077 m, 2041 s, 2002 s. On addition of 10 equiv of β-cyclodextrin,
5022
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
well with shifts occurring primarily for the C3 and C5 protons of
the cyclodextrin and the aryl protons on the aromatic ring of the
model complex. Interestingly, for 3Na changes are also observed
for the methylene protons of the PTA ligand with the resonance
at 4.05 ppm shifting downfield to 4.09 ppm. While not large, this
shift could still indicate some interaction between the PTA ligand
and the cyclodextrin, possibly through hydrogen bonding be-
tween the hydroxyl groups of the cyclodextrin and the amines.
However, as only one set of signals is observed for both the host
and guest components, the strength of the binding or rate of
dissociation and association does not appear to have been
significantly altered, relative to 1Na, by any such interaction.
Complex 5Na causes significantly different changes in the
chemical shifts of the cyclodextrin proton resonances. As the ratio
of 5Na:CyD is increased the C5 proton resonance behaves in a
manner similar to the other complexes, shifting upfield from 3.75 to
3.64 ppm. However, the C3 proton resonance now shifts downfield
from 3.89 to 3.94 ppm. Additionally, there is an upfield shift for the
C4 proton resonance from 3.51 to 3.48 ppm. Despite the difference
in the resonances for the cyclodextrin, the aryl protons on the diiron
complex show a similar upfield shift upon increasing the ratio of
cyclodextrin to 5Na as observed for the other complexes. All of
these changes indicate that 5Na interacts with cyclodextrins in
solution; however, the discrepancies between the NMR data for
5Na and the other sulfonated complexes implies thatit interacts in a
different manner. One possibility is that the aryl sulfonate is
entering the cyclodextrin cavity from the primary hydroxyl side,
preventing full penetration and positioning the sulfonate group
closer to the C3 proton.
Electrochemical Studies of 2Naꢀ5Na in MeCN. In MeCN,
the complexes show similar electrochemical events as have been
reported for the analogous diiron systems with different
dithiolates,5 with both an irreversible FeIFeI/FeIFe0 reduction
and irreversible FeIFeI/FeIFeII oxidation, Figure 4 and Table 5.
For complexes 2Na and 3Na, a second oxidative event appears at
about 100 mV. This event is not present when the reductive scan
is stopped prior to the FeIFeI/FeIFe0 reduction and is thus
attributed to a degradation product of the first reduction. While
all of the oxidative events follow the expected trend based on the
donor abilities of the ligands as established from ν(CO) infrared
studies, the reductive events do not, and instead the monosub-
stituted complexes are all reduced at ∼ꢀ1.8 V. More interesting
is the fact that the stronger donor containing complex, 3Na, is
reduced at slightly more positive potentials than either 2Na or
4Na. The differences are more pronounced in their electroche-
mical response to HOAc, added as a proton source for electro-
catalytic reduction of Hþ to H2.
Table 5. Electrochemical Data for 2Naꢀ5Na in MeCNa
complex
Epc (V)
Epa (V)
1Na12
2Na
ꢀ1.61
ꢀ1.85
ꢀ1.81
ꢀ1.85
ꢀ2.24
0.59
0.40
3Na
0.22
4Na.
5Na
0.31
ꢀ0.19
a Ar deaerated CH3CN solution (0.1 M Bu4NBF4). All experiments
were recorded using a glassy carbon working electrode (A = 0.071 cm2)
referenced to Cp2Fe/Cp2Feþ as an internal standard and a Pt counter
electrode at a scan rate of 200 mV/s.
n
For complex 2Na dissolved in MeCN, upon addition of HOAc
an increase in the current response of the reductive event
Figure 5. Cyclic voltammograms of complexes 2Na, 3Na, 4Na, and 5Na in MeCN with 0.1 M Bu4NBF4 at a scan rate of 200 mV/s showing the effect of
added HOAc. Concentration of all samples is 2 mM.
5023
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Table 6. Electrochemical Data for 2Naꢀ5Na in H2Oa
Complex
Epc (V) [With CyD]
Epa (V) [With CyD]
1Na12
2Na
ꢀ1.22 [ꢀ1.31]
ꢀ1.45 [ꢀ1.47]
ꢀ1.39 [b]
b
0.78 [0.82]
0.71 [0.68]
0.54 [0.53]
0.13 [0.13]
3Na
5Na
a Ar deaerated H2O solution (0.01 M NaCl). All experiments were
recorded using a glassy carbon working electrode (A = 0.071 cm2)
referenced to Ag/AgCl (sat’d KCl) and a Pt counter electrode at a scan
rate of 200 mV/s. b Reduction is beyond the solvent window.
substituted complexes 5Na shows a response to HOAc at the first
reduction potential, that with subsequent additions of acid splits
into two distinguishable redox events at ꢀ2.15 and ꢀ2.26 mV.
Figure 7 shows plots of the change in current versus the
equivalents of acetic acid. The slopes for best fit lines in all of
the complexes indicate a similar response for the reduction of Hþ
from HOAc.7
Electrochemical Studies of 2Naꢀ5Na in H2O. The increased
water solubility of the complexes permits electrochemical examina-
tion in aqueous solutions. Similar to MeCN solutions, most of the
compounds in H2O show both an irreversible oxidation and
reduction, Figure 6 and Table 6. Similar to what was observed in
MeCN the reductive events of 2Na and 4Na are in a similar range,
ꢀ1.39 to ꢀ1.45 V, and are not entirely consistent with expected
changes resulting from the donor abilities of the ligands. For
complex 5Na the reductive event seen in MeCN solution is not
observed and is likely beyond the solvent window. While the data
for comparison of the aqueous electrochemistry to other diiron
model complexes is still strongly limited by the scarcity of water-
soluble systems, comparison of 3Na to the similar ligand substi-
tuted DAPTA complex by Sun and co-workers shows that the two
compounds have similar reductive events at ꢀ1.45 and ꢀ1.5
respectively.9 Complexes 2Na and 5Na adsorb onto the surface
of the glassy carbon electrode, resulting in changes in the observed
potentials for the redox events, and require continual cleaning of
the electrode to reproduce the events of the initial scan. At no point
do the sulfonate species react with water; they are surprisingly
stable, showing no change in the electrochemical response after
being left in deaerated solution for several hours.
In the presence of β-cyclodextrin, all of the complexes show
noted decreases in the current intensity associated with both the
reductive and the oxidative events. This is commonly seen in other
electrochemical studies of inclusion complexes and is attributed to
the slower diffusion of the hostꢀguest complex relative to the free
guest.36 The reductive events also show minor differences in the
potentials with the initial reduction of 3Na shifting more cathodic
by greater than 100 mV. (This is best seen by DPV displays.) By
comparison, the oxidative events only shift at most between 30
and 40 mV, also more cathodic. A change in the reduction event of
∼80 mV was noted for 1Na in the presence of cyclodextrin.12 The
hydrophobic nature of the cyclodextrin cavity likely plays a large
role in this difference with the neutral oxidized species being less
disfavored than the dianionic reduced species and thus shifted to
more accessible potentials while the reduction is shifted to more
negative potentials.15
Figure 6. Cyclic voltammograms of complexes (Top) 2Na; (Middle)
3Na; (Bottom) 5Na in 0.01 M aqueous NaCl in the absence (black line)
and presence (red line) of 1 mM β-cyclodextrin. Scans were run at 200
mV/s with sample concentration of 0.1 mM.
around ꢀ2.10 V is observed and distinguishable from the event
attributed to reduction of HOAc on glassy carbon, Figure 5. The
fact that this response occurs at potentials past the initial
reduction of 2Na is consistent with trends reported for the less
electron rich model complexes.17,30ꢀ35 Similarly, complexes 3Na
and 4Na exhibit an increase in the current well beyond their first
reductive events, ꢀ2.24 and ꢀ2.15 V, respectively. While both
2Na and 3Na show an anodic shift for this event at higher
concentrations of HOAc, the increase observed in 4Na shifts
to more positive potentials with incremental addition of the
acid similar to the behavior observed in 5Na. Unlike the other
The electrochemical response to HOAc was also studied for the
water-soluble complexes in aqueous solutions. Because the first
reduction is beyond the solvent window and complex 5Na does not
react with HOAc, the changes in its voltammograms with added
5024
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
Figure 7. Aqueous cyclic voltammograms of complexes 2Na in the absence (top, 7a) and presence (bottom, 7b) of 10 equiv of β-CyD, showing the
response to incremental addition of HOAc. Scans were run at 200 mV/s with samples concentration of 0.1 mM. Next to each set of voltammograms, 7a0
and 7b0, are plots of the change in current versus equivalents of HOAc for points near the two observed events ꢀ1450 (diamonds) and ꢀ1550 mV
(squares). Note that whereas the responses of the events in 7a/7a0 parallel each other, for 7b/7b0 the more positive event at ꢀ1450 mV is less responsive
in the presence of β-CyD.
HOAc are indistinguishable from voltammograms of pure HOAc
increments itself. However, for complexes 2Na and 3Na there are
notable changes in the current intensities for the reductive events
similar to that observed for the parent complex 1Na. Upon addition
of HOAc to a solution of 2Na, there is an increase in current for the
FeIFeI/FeIFe0 reduction at ꢀ1.45 mV, Figure 7. A second event
about 100 mV more cathodic is also observed to increase with
increasing [Hþ], and with a similar response slope, Figure 7a0. The
presence of multiple events for proton reduction is not surprising as
this is generally the case when proton reduction studies are carried
out in MeCN. Evans and co-workers have proposed that these
separate events, observed in MeCN for (μ-pdt)[Fe(CO)3]2, are
due to conformational changes that occur upon reduction resulting
in multiple conformersthat reduce protons at different potentials.37
Of further interest to this notion is that when this proton reduction
assay is carried out in the presence of β-cyclodextrin, Figure 7b, the
more anodic of the two responses is less affected by the β-CyD.
That this change in the ratio of these responses could be resulting
from limited mobility inside of the cyclodextrin cavity is an
intriguing prospect that parallels the concept that the geometry
of the enzyme active site can be held in place by the surrounding
peptide architecture.
derivatives has resulted in greatly increased water solubility as
compared to the widely studied propane-dithiolate analogues. The
aryl sulfonate approach shows significant advantages over reported
methods for developing water-soluble diiron complexes including
(1) the facile synthesis of a variety of water-soluble, CO-substituted
derivatives; (2) the stability of resulting complexes in oxygen free
water for several hours; and (3) the promotion of the interaction
between the diiron complex and β-cyclodextrin which greatly
enhances the water solubility.
Despite the advantages, this series has other problems; the
amphiphilic nature of the sulfonate complexes leads to adsorp-
tion of the compounds onto the glassy carbon electrode and
fouling of the electrode surface results. This problem is slightly
resolved by balancing out the complex using the more hydro-
philic PTA ligand or encapsulation in the cyclodextrins; however,
it is not completely eliminated. A more dynamic electrode
surface, such as that obtained through use of a mercury drop
electrode,7 could lead to a better understanding of the electro-
chemistry of these water-soluble species. It is clear that mechan-
istic studies of electrocatalytic behavior are best performed in
nonaqueous solvents which might in fact more accurately match
the dielectric of the [FeFe]-H2ase active site. In either case, once
optimal biomimetics are designed, through correlation of syn-
thetic advances, electrochemical, and mechanistic advances, the
engineering of solid phase matrixes with H2O delivery routes will
be required for applications.
’ CONCLUSIONS
The facile incorporation of the sulfonate group into the dithio-
late of the diiron hexacarbonyl and P-donor ligand substituted
5025
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026

Inorganic Chemistry
ARTICLE
’ ASSOCIATED CONTENT
(27) Singleton, M. L.; Jenkins, R. M.; Klemashevich, C. L.; Darensbourg,
M. Y. C. R. Chim. 2008, 11, 861.
(28) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y.
Chem. Rev. 2007, 107, 4273.
(29) Schneider, H.-J.; Hacket, F.; Ruediger, V.; Ikeda, H. Chem. Rev.
1998, 98, 1755.
S
Supporting Information. Full structure files, cyclic
b
voltammograms, and spectral data. This material is available free
of charge via the Internet at http://pubs.acs.org.
(30) Capon, J.-F.; Ezzaher, S.; Gloaguen, F.; Pꢀetillon, F. Y.;
Schollhammer, P.; Talarmin, J.; Davin, T. J.; McGrady, J. E.; Muir,
K. W. New J. Chem. 2007, 31, 2052.
(31) Gloaguen, F.; Morvan, D.; Capon, J.-F.; Schollhammer, P.;
Talarmin, J. J. Electroanal. Chem. 2007, 603, 15.
’ AUTHOR INFORMATION
Corresponding Author
*E-mail: marcetta@mail.chem.tamu.edu.
(32) Capon, J.-F.; Ezzaher, S.; Gloaguen, F.; Pꢀetillon, F. Y.;
Schollhammer, P.; Talarmin, J. Chem.—Eur. J. 2008, 14, 1954.
(33) Capon, J.-F.; Gloaguen, F.; Pꢀetillon, F. Y.; Schollhammer, P.;
Talarmin, J. Coord. Chem. Rev. 2009, 253, 1476.
(34) Chong, D.; Georgakaki, I. P.; Mejia-Rodriguez, R.; Sanabria-
Chinchilla, J.; Soriaga, M. P.; Darensbourg, M. Y. Dalton Trans.
2003, 4158.
(35) Tye, J. W.; Lee, J.; Wang, H.-W.; Mejia-Rodriguez, R.; Reibenspies,
J. H.; Hall, M. B.; Darensbourg, M. Y. Inorg. Chem. 2005, 44, 5550.
(36) Matsue, T.; Evans, D. H.; Osa, T.; Kobayashi, N. J. Am. Chem.
Soc. 1985, 107, 3411.
(37) Felton, G. A. N.; Petro, B. J.; Glass, R. S.; Lichtenberger, D. L.;
Evans, D. H. J. Am. Chem. Soc. 2009, 131, 11290.
’ ACKNOWLEDGMENT
We gratefully acknowledge financial support from the National
Science Foundation (Grant CHE-0910679) and the Robert
A. Welch Foundation (Grant A-0924). We appreciate the help
of Jason Denny in the final preparation of the structure tables
and figures.
’ REFERENCES
(1) Nicolet, Y.; Lemon, B. J.; Fontecilla-Camps, J. C.; Peters, J. W.
Trends Biochem. Sci. 2000, 25, 138.
(2) Hambourger, M.; Gervaldo, M.; Svedruzic, D.; King, P. W.; Gust,
D.; Ghirardi, M.; Moore, A. L.; Moore, T. A. J. Am. Chem. Soc. 2008,
130, 2015.
(3) Esswein, A. J.; Nocera, D. G. Chem. Rev. 2007, 107, 4022.
(4) Tard, C.; Pickett, C. J. Chem. Rev. 2009, 109, 2245.
(5) Felton, G. A. N.; Mebi, C. A.; Petro, B. J.; Vannucci, A. K.; Evans,
D. H.; Glass, R. S.; Lichtenberger, D. L. J. Organomet. Chem. 2009,
694, 2681.
’ NOTE ADDED AFTER ASAP PUBLICATION
This paper was published on the Web on April 27, 2011. The
in-text citations of refs 23 and 28 have been corrected and the
updated version was reposted on May 2, 2011.
(6) Le Cloirec, A.; Davies, S. C.; Evans, D. J.; Hughes, D. L.; Pickett,
C. J.; Best, S. P.; Borg, S. Chem. Commun. 1999, 2285.
(7) Felton, G. A. N.; Glass, R. S.; Lichtenberger, D. L.; Evans, D. H.
Inorg. Chem. 2006, 45, 9181.
(8) Mejia-Rodriguez, R.; Chong, D.; Reibenspies, J. H.; Soriaga,
M. P.; Darensbourg, M. Y. J. Am. Chem. Soc. 2004, 126, 12004.
(9) Na, Y.; Wang, M.; Jin, K.; Zhang, R.; Sun, L. J. Organomet. Chem.
2006, 691, 5045.
(10) Li, P.; Wang, M.; Chen, L.; Wang, N.; Zhang, T.; Sun, L.
CrystEngComm 2008, 10, 267.
(11) Apfel, U.-P.; Halpin, Y.; Gottschaldt, M.; Goerls, H.; Vos, J. G.;
Weigand, W. Eur. J. Inorg. Chem. 2008, 5112.
(12) Singleton, M. L.; Reibenspies, J. H.; Darensbourg, M. Y. J. Am.
Chem. Soc. 2010, 132, 8870.
(13) Breslow, R.; Dong, S. D. Chem. Rev. 1998, 98, 1997.
(14) Hapiot, F.; Tilloy, S.; Monflier, E. Chem. Rev. 2006, 106, 767.
(15) Johnson, M. D.; Bernard, J. G. Chem. Commun. 1996, 2, 185.
(16) Kano, K. Colloid Polym. Sci. 2008, 286, 79.
(17) Surawatanawong, P.; Tye, J. W.; Darensbourg, M. Y.; Hall,
M. B. Dalton Trans. 2010, 39, 3093.
(18) Spek, A. L. J. Appl. Crystallogr. 2003, 36, 7–13.
(19) (a) APEX2; Bruker AXS Inc.: Madison, WI, 2007. (b) GADDS;
Bruker AXS Inc.: Madison, WI, 2007. (c) SAINT; Bruker AXS Inc.:
Madison, WI, 2007. (d) Sheldrick, G. M. Acta Crystallogr. 2008,
A64, 112.
(20) Barbour, L. J. J. Supramol. Chem. 2003, 1, 189.
(21) Li, H.; Rauchfuss, T. B. J. Am. Chem. Soc. 2002, 124, 726.
(22) Lyon, E. J.; Georgakaki, I. P.; Reibenspies, J. H.; Darensbourg,
M. Y. J. Am. Chem. Soc. 2001, 123, 3268.
(23) Lyon, E. J.; Georgakaki, I. P.; Reibenspies, J. H.; Darensbourg,
M. Y. Angew. Chem. 1999, 38, 3178.
(24) Liu, Y.-C.; Tu, L.-K.; Yen, T.-H.; Lee, G.-H.; Chiang, M.-H.
Dalton Trans. 2011, 40, 2528.
(25) Wei, C. H.; Dahl, L. F. Inorg. Chem. 1965, 4, 1.
(26) Li, P.; Wang, M.; He, C.; Li, G.; Liu, X.; Chen, C.; Åkermark, B.;
Sun, L. Eur. J. Inorg. Chem. 2005, 2506.
5026
dx.doi.org/10.1021/ic200272x |Inorg. Chem. 2011, 50, 5015–5026