A convenient procedure for the synthesis of fluoro-iron(III) complexes of common synthetic porphyrinates

Article abstract of DOI:10.1142/S108842461450014X

We report here an improved method for the preparation of fluoro-iron(III) porphyrinate complexes. Treatment of [Fe₂(P)₂(μ-O)] (P = tetraphenylporphyrinate {TPP}, tetra-p-tolylporphyrinate {TTP}, or octaethylporphyrinate {OEP}) or [Fe(OH)(OH₂)(TMP)] (TMP = tetramesitylporphyrinate) with the commercially available fluorinating agent, Et₃N·3HF, in dichloromethane affords the desired [FeF(P)] complexes in a straightforward fashion and in good yield while avoiding the use of aqueous hydrofluoric acid. All fluoro-iron(III) complexes have been completely characterized by a series of different spectroscopic techniques including cyclic voltammetry. Reaction of a representative complex, [FeF(OEP)], with various chloride reagents demonstrates that halide exchange with chloride is facile, but only proceeds at an appreciable rate in the presence of proton sources. Unexpectedly, treatment of [FeF(OEP)] with NOBF₄ did not to lead formation of an oxidized species, but rather to formation of the {Fe-NO}⁶ complex, [Fe(NO)(OEP)](BF₄).

Full text of DOI:10.1142/S108842461450014X

Journal of Porphyrins and Phthalocyanines
J. Porphyrins Phthalocyanines 2014; 18: 416–423
DOI: 10.1142/S108842461450014X
Published at http://www.worldscinet.com/jpp/
A convenient procedure for the synthesis of fluoro-iron(III)
complexes of common synthetic porphyrinates
Daniel J. Meininger, Nicanor Muzquiz, Hadi D. Arman and Zachary J.Tonzetich*◊
Department of Chemistry, University of Texas at San Antonio (UTSA), One UTSA Circle, San Antonio,
TX 78249, USA
Received 29 January 2014
Accepted 4 March 2014
ABSTRACT: We report here an improved method for the preparation of fluoro-iron(III) porphyrinate
complexes. Treatment of [Fe (P) (m-O)] (P = tetraphenylporphyrinate {TPP}, tetra-p-tolylporphyrinate
2
2
{
TTP}, or octaethylporphyrinate {OEP}) or [Fe(OH)(OH )(TMP)] (TMP = tetramesitylporphyrinate)
2
with the commercially available fluorinating agent, Et N·3HF, in dichloromethane affords the desired
3
[
FeF(P)] complexes in a straightforward fashion and in good yield while avoiding the use of aqueous
hydrofluoric acid. All fluoro-iron(III) complexes have been completely characterized by a series of
different spectroscopic techniques including cyclic voltammetry. Reaction of a representative complex,
[
FeF(OEP)], with various chloride reagents demonstrates that halide exchange with chloride is facile,
but only proceeds at an appreciable rate in the presence of proton sources. Unexpectedly, treatment of
[
{
FeF(OEP)] with NOBF did not to lead formation of an oxidized species, but rather to formation of the
4
6
Fe–NO} complex, [Fe(NO)(OEP)](BF ).
4
KEYWORDS: iron(III) porphyrinates, fluoride complexes, synthetic procedures.
In the area of porphyrin chemistry, examples of
INTRODUCTION
terminal fluoride and difluoride complexes have been
reported for several transition metals [8–21]. Recently,
a high-valent Mn porphyrin fluoride complex was
demonstrated to carry out catalytic fluorination with
excellent efficiency [22, 23]. Among transition metal
porphyrinates, however, the most well-studied fluoride
complexes are those of iron(III). Fluoro-iron(III) por-
phyrinates have been structurally characterized for a
series of different porphyrin ligands and various aspects
of their spectroscopy have been examined in detail
Examples of transition metal complexes containing
fluoride ligands remain relatively scarce in comparison
to those containing the heavier halogens [1–5]. Despite
a recent surge in interest due to their ability to effect
catalytic C–F bond forming reactions [6], metal fluoride
complexes are often neglected in studies of metal halide
compounds because of their difficulty of preparation. In
contrast to binary metal chlorides, bromides, and iodides,
binary metal fluorides (MF ) are often unavailable,
x
especially if lower metal oxidation states are desired
[
24–32]. Particular attention has been paid to the effects
[
7]. Introduction of fluoride ligands is therefore usually
of fluoride ligation in modulating the reactivity of high
valent oxo-iron porphyrinates and in determining the
electronic structure of porphyrin p-cation radicals [33–44].
In addition, fluoro-iron(III) porphyrinates are also of
interest as models for fluoride interaction with heme
proteins [45–48].
accomplished at a later stage in the metal complex
synthesis by halogen exchange. Halogen exchange is
not always straightforward, however, typically requiring
large excesses of fluoride salts such as NaF or KF, or
hydrofluoric acid.
During the course of recent investigations into the
chemistry of iron(III) porphyrinates containing silane-
thiolate ligands, we discovered that treatment of oxygen-
ligated iron(III) porphyrinates with the easy to use,
◊SPP full member in good standing
*
Correspondence to: Zachary J. Tonzetich, tel: +1 (210)-458-5465,
email: zachary.tonzetich@utsa.edu
Copyright © 2014 World Scientific Publishing Company

A CONVENIENT PROCEDURE FOR THE SYNTHESIS OF FLUORO-IRON(III) COMPLEXES
417
commercially available HF source, Et N·3HF, resulted in
nearquantitativeformationoffluoro-iron(III)porphyrinates
(116–167 mmol Fe) of the desired oxo-bridged iron(III)
porphyrinate, [Fe (P) (m-O)], or hydroxide complex,
3
2
2
[49]. We have now optimized this protocol and found it
[Fe(OH)(H O)(TMP)]. The solid is dissolved completely
2
to be general for a variety of synthetic porphyrin ligands.
We report here this new fluorination procedure along with
detailed characterization of the fluoride complexes.
in 25 mL of CH Cl and 40.0 mL (245 mmol) of Et N·3HF
2
2
3
is added. The resulting mixture is allowed to stir for 18 h at
ambient temperature (20–23°C). The solvent is removed
in vacuo and the resulting purple microcrystalline solid
is washed with cold acetone yielding the desired fluoride
complex. Yields and characterization data for specific
porphyrinate ligands appear below. Spectroscopic data
EXPERIMENTAL
(
UV-vis, IR, and EPR) matched those available in the
General comments
literature [24, 25, 28, 38, 53, 54]. All fluoro complexes
were found to retain variable amounts of water, which
could be observed as a broad singlet near 0 ppm in
the H NMR spectrum. In the case of the TPP and
TTP complexes, insufficient solubility in benzene-d6
precluded definitive assignment of the ortho-aryl protons
Unless noted, manipulations were performed under
ambient atmosphere. Tetrahydrofuran, diethyl ether,
methylene chloride, pentane, and toluene were purified
by sparging with argon and passage through two columns
packed with 4 Å molecular sieves. Benzene, benzene-d6,
and 2-methyltetrahydrofuan were dried over sodium
1
1
by H NMR.
1
then vacuum-distilled. H NMR spectra were recorded in
1
[
FeF(TPP)]. 87% yield. H NMR: d , ppm 81.2 (v br
H
benzene-d on a Varian INOVA spectrometer operating at
6
s, 8 pyr-CH), 11.0 (br s, 4 m-ArH), 10.2 (br s, 4 m-ArH),
6
3
5
00 MHz and referenced to the residual C D H peak of
-1
-1
6
5
.67 (s, 4 p-ArH). UV-vis (toluene): lmax,nm (e, M .cm )
the solvent (d 7.16 ppm vs. TMS). UV-vis spectra were
recorded at ambient temperature in toluene on a Cary-
28 (34,000), 415 (140,000), 534 (sh), 594 (6700), 634
-1
(
sh). IR (KBr): n, cm 3054 (w), 1597 (m), 1485 (m),
6
0 spectrophotometer in Teflon-capped quartz cells.
1
440 (m), 1340 (m), 1202 (m), 1175 (m), 1071 (m), 1004
vs), 996 (s), 804 (s), 750 (m), 720 (s), 702 (m), 660 (m),
11 (m, nFe-F). EPR: g = 5.77, g = 1.98. Anal. calcd. for
Cyclic voltammetry measurements were performed in
dichloromethane in a single compartment cell under a
nitrogen atmosphere (in the glovebox) at 25°C using a
CH Instruments 620D electrochemical workstation. A
(
6
⊥
||
C H FFeN ·1.5H O: C, 73.95; H, 4.37; N, 7.84. Found
4
4
28
4
2
C, 73.85; H, 3.94; N, 7.92.
FeF(TTP)]. 89% yield. H NMR: d , ppm 81.0 (v
br s, 8 pyr-CH), 10.95 (s, 4 m-ArH), 10.17 (s, 4 m-ArH),
3
-electrode set-up was employed comprising a 1 mm
1
[
H
diameter Pt disk working electrode, platinum wire
auxiliary electrode, and Ag quasi-reference electrode.
Triply recrystallized Bu NPF or Bu NBF was used as
the supporting electrolyte. All electrochemical data were
referenced to the ferrocene/ferrocenium couple at 0.00 V.
EPR measurements were recorded in 4 mm quartz tubes
on a Bruker E500 EPR spectrometer operating at the
X-band (9.33 GHz) at a modulation frequency of 100 kHz
and modulation amplitude of 10G. Low temperature
measurements were made in frozen 2-MeTHF glasses at
-1
6
.68 (s, 12 p-ArCH ). UV-vis (toluene): l , nm (e, M .
3
max
-1
4
6
4
4
cm ) 329 (28,000), 415 (120,000), 590 (5200), 635 (sh).
-1
IR (KBr): n, cm 3023 (w), 1596 (m), 1485 (m), 1440
(m), 1334 (m), 1201 (m), 1174 (m), 1070 (m), 1003 (vs),
9
95 (s), 805 (s), 749 (s), 721 (m), 702 (s), 660 (m), 612
(
m, nFe-F). EPR: g = 5.77, g = 1.98. Anal. calcd. for
⊥
||
C H FFeN ·0.5H O: C, 76.60; H, 4.95; N, 7.44. Found
4
8
38
4
2
C, 76.38; H, 4.93; N, 7.47.
FeF(OEP)]. 85% yield. H NMR: d , ppm 34.9 (br
s, 8 Et-CH ), 34.2 (br s, 8 Et-CH ), 5.10 (s, 24 Et-CH ),
1
[
H
8
0 K with temperature control maintained by a nitrogen
2
2
3
(
77K) flow cryostat (ESR900, Oxford Instruments,
-
(
5
(
1
34.0 (v br s, 4 meso-CH). UV-vis (toluene): lmax, nm
-1 -1
Inc.). Elemental analyses were performed by Atlantic
Microlab, Inc. of Norcross, GA.
e, M .cm ) 347 (47,000), 394 (130,000), 478 (12,000),
-1
06 (sh), 591 (11,000). IR (KBr): n, cm 2964 (vs), 2930
s), 2868 (s), 1466 (m), 1445 (m), 137 (m), 1315 (m),
267 (m), 1214 (m), 1146 (s), 1111 (m), 1055 (s), 1014
s), 981 (m), 957 (s), 915 (m), 843 (m), 747 (m), 734 (m),
Materials
(
Metalloporphyrinates, [Fe (TPP) (m-O)], [Fe (TTP) -
2
2
2
2
7
1
7
19 (m), 703 (m), 600 (m, nFe-F). EPR: g = 5.83, g =
(
m-O)], [Fe (OEP) (m-O)], and [Fe(OH)(H O)(TMP)],
⊥ ||
2
2
2
.99. Anal. calcd. for C H FFeN ·H O: C, 69.11; H,
were prepared by literature procedures [50–52]. Et N·3HF
36 44
4
2
3
.41; N, 8.93. Found C, 69.32; H, 7.18; N, 8.98.
and NOBF were purchased from commercial suppliers
4
1
[
FeF(TMP)]. 81% yield. H NMR: d , ppm 80.1 (v br
(
Aldrich and Strem, respectively) and used as received.
H
s, 8 pyr-CH), 12.35 (s, 4 m-ArH), 11.49 (s, 4 m-ArH), 4.5
v br s, 12 o-ArCH ), 3.32 (s, 12 p-ArCH ), 2.9 (v br s, 12
Et NHCl was prepared by treatment of Et N with
3
3
(
HCl·dioxane (4.0 M) in diethyl ether.
3
3
-1
-1
o-ArCH ). UV-vis (toluene): l , nm (e, M .cm ) 323
3
max
(36,000), 416 (130,000), 593 (5900), 636 (sh). IR (KBr):
Synthesis
-1
n, cm 3020 (m), 2917 (m), 2870 (m), 1610 (m), 1478
(m), 1440 (m), 1387 (m), 1328 (m), 1202 (m), 1063 (m),
1000 (vs), 868 (m), 852 (m), 832 (s), 804 (s), 725 (m),
General procedure for preparation of [FeF(P)].
A 50 mL round bottom flask is charged with 100 mg
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 417–423

418
D.J. MEININGER ET AL.
6
14 (m, nFe-F). EPR: g = 5.79, g = 1.98. Anal. calcd. for
axis along b explains the partially fulfilled systematic
absences. Inspection of the systematic absence statistics
clearly indicated the presence of a glide plane along
c. However, the twofold screw axis was unclear, with
65% of the 32 reflections that should be absent for the
⊥
||
C H FFeN ·0.5H O: C, 77.77; H, 6.18; N, 6.46. Found
C, 77.41; H, 6.15; N, 6.49.
5
6
52
4
2
Preparation of [Fe(NO)(OEP)](BF ) from [FeF­
4
(
OEP)]. A vial was charged with 30.1 mg (43.8 µmol)
of [FeF(OEP)] and 7.4 mg (63 µmol) of NOBF . The
monoclinic 2 axis being observed but at significantly
4
1
solids were suspended in 4 mL of CH Cl and allowed
weaker intensity (by a factor of 3) than the rest of the
data.
2
2
to stir for 18 h at ambient temperature in the glovebox
under an atmosphere of nitrogen. All volatiles were
removed in vacuo and the resulting purple solid was
triturated with pentane and collected by filtration. Mass:
RESULTS AND DISCUSSION
2
0.0 mg (58% yield). Spectroscopic features were found
The standard procedure for synthesis of fluoro-iron(III)
porphyrinates makes use of aqueous hydrofluoric acid
as a fluorinating agent [28, 29]. In this procedure, the
oxo-bridged iron(III) porphyrinate, [Fe (P) (m-O)], is
treated with HF (aq) to afford the desired fluoro-iron(III)
complex, [FeF(P)]. Low yields, the need to avoid glass
reaction vessels, and the danger associated with handling
HF (aq) render this procedure un-ideal as a practical
means of preparing the fluoride complexes of iron(III)
porphyrinates. Other procedures have also been reported,
including direct metallation of the porphyrin ligand with
+
to match those of [Fe(NO)(OEP)] in the literature [55].
1
H NMR (CD Cl ): d , ppm 7.96 (br s, 4 meso-CH),
2
2
H
3
.14 (br s, 16 Et-CH ), 1.41 (br s, 24 Et-CH ). UV-vis
2
3
2
2
-
1
(
1
CH Cl ): lmax, nm 361, 383 (sh), 555. IR (KBr): n, cm
2
2
838 (nNO).
X­ray crystallography
A crystal of [FeF(TTP)]·2CH Cl suitable for X-ray
diffraction was mounted in Paratone oil onto a glass
fiber and frozen under a nitrogen cold stream maintained
by an X-Stream low-temperature apparatus. The data
were collected at 98(2) K using a Rigaku AFC12/Saturn
2
2
FeF , although the generality of such procedures has not
3
been established [27]. We therefore sought to examine
the utility of a new fluorination protocol involving
treatment of oxygen-ligated iron(III) porphyrinates with
7
24 CCD fitted with Mo Ka radiation (l = 0.71073 Å).
Data collection and unit cell refinement were performed
using Crystal Clear software [56]. The total number of
data were measured in the range 3.1 < q < 27.6° using w
scans. Data processing and absorption correction, giving
minimum and maximum transmission factors, were
accomplished with Crystal Clear and ABSCOR [57],
respectively. The structure was solved by direct methods
Et N·3HF. We had identified this reaction previously
3
during studies of iron(III) porphyrinates containing
silanethiolate ligands [49]. At that time, we observed that
the hydroxo-iron(III) complex, [Fe(OH)(H
reacted readily with Et N·3HF to afford the fluoro-
O)(TMP)],
2
3
iron(III) porphyrinate, [FeF(TMP)] (Equation 1). We
reasoned that related oxo-bridged iron(III) porphyrinates
containing ligands other than TMP might react in similar
2
and refined on F using full-matrix, least-squares
techniques with SHELXL-97 [58, 59]. Non-hydrogen
atoms were refined with anisotropic displacement
parameters. All carbon bound hydrogen atom positions
were determined by geometry and refined by a riding
model.
fashion with Et
synthesis for a series of different fluoro-iron(III) species
[53]. Et N·3HF is a commercially available fluorinating
agent that can be handled in straightforward fashion
with standard laboratory protocol thereby obviating the
need for plastic reaction vessels. Previous work has also
demonstrated the utility of this reagent in converting
oxygen bound ligands to fluorides with other transition
metal platforms [60].
3
N·3HF thereby providing a convenient
3
The crystal of [FeF(TTP)]·2CH Cl was found to be
2
2
twinned, displaying a two-fold axis. The TWIN command
was used to account for this condition during refinement.
Due to the severe disorder in the co-crystallized CH Cl
2
2
molecules, the PLATON SQUEEZE command was used
to remove one of the two molecules, resulting in a 132 Å
Reactionsofaseriesofoxo-bridgediron(III)complexes
containing porphyrin ligands featuring meso and pyrrolic
3
void. The structure also exhibits pseudo-symmetry. The
substitution with Et N·3HF in dichloromethane were
3
correct space group should be P2 /c, however, Pc gave
a better description of the structure. The fact that the
structure shows almost but not quite a twofold screw
found to afford the desired fluoride complexes in excellent
yield (Scheme 1). Confirmation of the presence of the
1
fluoride ligand was provided by IR spectroscopy (nFe-F
)
Mes
OH
Mes
N
F
N
Mes
CH2Cl2
N
Mes
N
Fe
+
Et3N·3HF
Fe
(
1)
N
Mes
N
N
Mes
N
OH2
Mes
Mes
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 418–423

A CONVENIENT PROCEDURE FOR THE SYNTHESIS OF FLUORO-IRON(III) COMPLEXES
419
R
by slow evaporation of a solution of
N
[
FeF(TTP)]·xH OinCH Cl (Fig.1).This
2 2 2
N
N
Fe
F
complex is the only member of the series
of fluoride complexes described here not
disclosed previously. Similar to reported
fluoro-iron(III) porphyrinate structures,
the solid-state structure of [FeF(TTP)]
displays a short Fe–F bond distance
of 1.805(4) Å [28, 62]. The complex
features two molecules of CH Cl in
N
CH2Cl2
N
O
+
Et3N·3HF
N
Fe
N
N
N
N
Fe
R
N
N
for TPP, TTP, and OEP
R
Scheme 1. Synthesis of fluoro-iron(III) porphyrinates
2
2
the asymmetric unit. Unfortunately, a
crystallographic disorder prevented satis-
and by comparison of the UV-vis spectra to previous
work (see Experimental section). A moderate excess
factory refinement of both solvent
molecules (see Experimental section and Supporting
information). Nonetheless, the positions of the CH Cl
2
of Et N·3HF was found to be optimal, as prolonged
3
2
exposure to CH Cl in the presence of proton sources was
molecules suggest a hydrogen bonding interaction with
the fluoride ligand is present in the solid state.
2
2
found to lead to formation of chloro-iron(III) complexes
vide infra). Washing the isolated fluoride complexes
(
Previous studies on fluoro-iron(III) porphyrinates
have examined many of their spectroscopic features
[40, 63–68], including their electrochemical behavior
[53, 54, 69]. For completeness, we have examined the
electrochemistry of each of the fluoride complexes
described above in dichloromethane. All species show
three reversible or quasi-reversible electrochemical
events by cyclic voltammetry in dichloromethane (Fig. 2,
Table 1, and see Supporting information). Interestingly,
reduction to iron(II) is nearly reversible in all cases,
whereas the analogous reduction event is completely
irreversible in the chloro-iron(III) complexes (see
Supporting information). In contrast to the first reduction
event, the first and second oxidation events for both chloro
and fluoro complexes are nearly identical, consistent with
these events corresponding to ligand-based processes
[36, 42, 46, 69–72].
with cold acetone was found to remove excess Et N·3HF
3
affording pure material. All attempted purification of
the complexes by alumina chromatography, however,
resulted in recovery of a mixture of [Fe (P) ]-m-O and
2
2
[
FeCl(P)].
In all cases, the fluoride complexes were found to
retain several equivalents of water, which could be
1
observed in the H NMR spectrum as a broad singlet near
0
ppm and accounted for during combustion analysis. We
believe the presence of water results from the desire of
the fluoride ligand to retain hydrogen-bond donors in the
secondary coordination sphere [40, 61]. This proposal is
corroborated by the fact that all known crystal structures
of fluoro-iron(III) porphyrinates feature co-crystallized
CHCl3 molecules engaging in a hydrogen-bonding
interaction with the fluorine atom [24, 26, 28]. Attempts
to grow crystals of the fluoro-iron(III) porphyrinates in
the present study in order to observe hydrogen-bonding
with water molecules were unsuccessful. However, the
crystal structure of [FeF(TTP)]·2CH Cl was obtained
The reversibility of the first oxidation wave in
[FeF(OEP)] prompted us to consider its chemical
oxidation in order to determine whether a fluoro-
iron porphyrin p-cation radical complex could
2
2
Fig. 1. Thermal ellipsoid (50%) rendering of the solid-state structure of [FeF(TTP)]·2CH Cl . Hydrogen atoms, minor components of
2
2
the disorder and co-crystallized CH Cl molecules omitted for clarity. Selected bond distances (Å): Fe(1)–F(1) = 1.805(4); Fe(1)–N(1) =
2
2
2
.052(4); Fe(1)–N(2) = 2.022(4); Fe(1)–N(3) = 2.076(4); Fe(1)–N(4) = 2.090(4); displacement of Fe(1) from N plane = 0.437 Å
4
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 419–423

420
D.J. MEININGER ET AL.
the likelihood for fluoride abstraction. Surprisingly,
treatment of [FeF(OEP)] with NOBF did not lead
4
to formation of [FeF(OEP)](BF ), but rather to the
4
oxidized nitrosyl complex, [Fe(NO)(OEP)](BF ),
4
as judged by NMR, IR, and UV-vis spectroscopy
(
Equation 2). At this time we have no data to suggest
a possible stoichiometry for this reaction or the fate
of the fluoride ligand. However, the procedure proved
reliable on multiple attempts with isolated yields
approaching 60%.
Given the facile nature of axial ligand exchange in
iron(III) porphyrinates, we next examined the stability
of the [FeF(P)] complexes in the presence of various
chloridesources. Prolongedexposure(days)ofthefluoro-
iron(III) complexes to dichloromethane or chloroform
was found to result in conversion to the chloro-iron(III)
complexes as judged by NMR spectroscopy. In solutions
of chloroform-d that had not been treated to remove
traces of HCl, this conversion was found to take place
within a matter of minutes. Conversely, reaction of a
representative complex, [FeF(OEP)], with Bu NCl was
4
Fig. 2. Electrochemical events observed for [FeF(OEP)] in
dichloromethane at a platinum electrode. Top trace is the
DPV and the bottom trace is the CV (50 mV/s scan rate). The
found to result in little to no conversion to the chloro-
iron complex as judged by UV-vis spectroscopy. This
result suggests that halide exchange at the iron(III)
porphyrinate is facilitated by proton sources, a
conclusion consistent with the observed instability of the
fluoro-iron complexes towards alumina chromatography.
To test this hypothesis more concretely, we examined
supporting electrolyte is 0.1 M Bu NPF6
4
Table 1. Electrochemical data for [FeF(P)] complexes prepared
‡
in this study
0/−
0/+
[FeF(P)]+/2+
Compound
[FeF(P)]
[FeF(P)]
the reaction of [FeF(OEP)] with Et NHCl. Addition of
3
[
[
[
FeF(TPP)]
FeF(TTP)]
FeF(OEP)]
FeF(TMP)]
-1.03
-1.01
-1.12
-1.21
+0.63
+0.66
+0.53
+0.62
+1.13
+1.13
+1.11
+1.10
a dichloromethane solution of Et NHCl to [FeF(OEP)]
lead to immediate conversion to the corresponding
chloro-iron complex (Equation 3). Conducting the same
3
reaction in pure toluene where Et NHCl is sparingly
3
[
soluble provided the opportunity to monitor this
conversion by UV-vis. Figure 3 displays the spectral
changes associated with the reaction of [FeF(OEP)] with
‡
Potentials represent E values determined by cyclic voltammetry
½
+
in CH Cl at a platinum electrode.All potentials reported vs. Fc/Fc .
2
2
Et NHCl. As is evident from the spectrum, conversion
3
of [FeF(OEP)] to [FeCl(OEP)] is clean displaying
several well defined isosbestic points. Thus, halide
exchange reactions of the fluoro-iron(III) porphyrinate
appear to require a proton source. To further confirm this
conclusion, we revisited the reaction of [FeF(OEP)] with
be isolated. The analogous complex containing a
chloride ligand bound to iron has been described for
a series of different porphyrin ligands [73–75]. As a
chemical oxidant we examined NOBF reasoning that
4
it would be sufficiently oxidizing while minimizing
Bu NCl in the presence of a proton source. Accordingly,
4
BF4
Et
F
Et
NO
Et
Et
Et
Et
CH2Cl2
N
N
Et
N
Fe
+
NOBF4
Et
N
Fe
(2)
N
Et
N
Et
N
N
Et
Et
Et
Et
Et
Et
Et
Et
F
Cl
Et
Et
Et
Et
N
N
Et
N
Fe
+
Et3NHCl
Et
N
Fe
(
3)
N
Et
N
Et
N
N
Et
Et
Et
Et
Et
Et
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 420–423

A CONVENIENT PROCEDURE FOR THE SYNTHESIS OF FLUORO-IRON(III) COMPLEXES
421
CONCLUSION
In this contribution we have described a facile new
procedure for the preparation of a series of fluoro-
iron(III) porphyrinates. This new method utilizes the
commercially available fluorinating agent, Et N·3HF, and
3
avoids the difficulties associated with handling aqueous
hydrofluoric acid. Purification of the complexes is facile
and isolated yields are greater than 80%. Studies with the
fluoro-iron complexes demonstrate the sensitivity of the
compounds towards halide exchange with chloride in the
presence of proton sources. Thus, an effort to minimize
prolonged exposure to chlorinated solvents should be
made when handling fluoro-iron(III) porphyrinates.
Acknowledgements
This work was supported by a grant from the Welch
Fig. 3. Electronic absorption spectra of the reaction of
FeF(OEP)] (­­­) with excess Et NHCl in toluene showing
conversion to [FeCl(OEP)] (—)
Foundation (AX-1772 to ZJT.). NM acknowledges the
ACS Project SEED program for financial support.
[
3
Supporting information
Additional spectra and the cif file for [FeF(TTP)]·
2
CH Cl (Figs S1–S13 and Table S1) are given in the
2 2
supplementary material. This material is available free of
charge via the Internet at http://www.worldscinet.com/
jpp/jpp.shtml.
Coordinates for [FeF(TTP)]·2CH Cl have been
2
2
deposited at the Cambridge Crystallographic Data
Centre (CCDC) under number CCDC-990086. Copies
can be obtained on request, free of charge, via www.
ccdc.cam.ac.uk/data_request/cif or from the Cambridge
Crystallographic Data Centre, 12 Union Road,
Cambridge CB2 1EZ, UK (fax: +44 1223-336-033 or
email: deposit@ccdc.cam.ac.uk).
REFERENCES
1
2
3
4
. Murphy EF, Murugavel R and Roesky HW. Chem.
Rev. 1997; 97: 3425–3468.
. Massa W and Babel D. Chem. Rev. 1988; 88:
Fig. 4. Electronic absorption spectra of [FeF(OEP)] and 1
equiv. Bu NCl in toluene/THF before (­­­) and after addition of
4
2
75–296.
. Fagnou K and Lautens M. Angew. Chem. Int. Ed.
002; 41: 26–47.
. Pagenkopf BL and Carreira EM. Chem.-Eur. J.
999; 5: 3437–3442.
p-TsOH (— 1 equiv., grey; — 2 equiv., black). Initial spectrum
is identical to that of [FeF(OEP)] and the final spectrum is
identical to that of [FeCl(OEP)]
2
1
addition of p-toluenesulfonic acid (p-TsOH) to a mixture
5. Winfield JM. J. Fluorine Chem. 1986; 33: 159–177.
6. Liang T, Neumann CN and Ritter T. Angew. Chem.
Int. Ed. 2013; 52: 8214–8264.
of [FeF(OEP)] and Bu NCl led to immediate conversion
4
to [FeCl(OEP)] as judged by UV-vis spectroscopy
(
Fig. 4). These results coupled with the apparent desire
7. Clark HCS andHolloway JH. In AdvancedInorganic
Fluorides, Nakajima T, Žemva B and Tressaud A.
(Eds.) Elsevier: Switzerland, 2000; pp 51–78.
8. Lecomte C, Protas J, Richard P, Barbe JM and Guilard
R. J. Chem. Soc., Dalton Trans. 1982: 247–250.
of the fluoro-iron(III) porphyrinates to associate with
hydrogen-bond donors (vide supra) strongly points to an
important role for hydrogen-bonding in mediating halide
exchange [25, 26].
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 421–423

422
D.J. MEININGER ET AL.
9
. Marchon JC, Latour JM, Grand A, Belakhovsky M,
Loos M and Goulon J. Inorg. Chem. 1990; 29: 57–67.
0. Okada K, SumidaA, Inagaki R and Inamo M. Inorg.
Chim. Acta 2012; 392: 473–477.
36. Dey A and Ghosh A. J. Am. Chem. Soc. 2002; 124:
3206–3207.
37. Nam W, Lim MH and Oh S-Y. Inorg. Chem. 2000;
39: 5572–5575.
1
1
1
1
1
1
1
1
1
1
2
2
2
2
2
2
2
2
2
2
3
3
3
3
1. La Mar GN and Walker FA. J. Amer. Chem. Soc.
38. Paeng K-J, Shin D-H and Paeng IR. Inorg. Chem.
1997; 36: 2901–2903.
39. Czarnecki K, Nimri S, Gross Z, Proniewicz LM
and Kincaid JR. J. Am. Chem. Soc. 1996; 118:
2929–2935.
40. Paeng IR and Paeng K-J. Bull. Korean Chem. Soc.
1995; 16: 997–1000.
41. Gross Z and Nimri S. Inorg. Chem. 1994; 33:
1731–1732.
42. Nanthakumar A and Goff HM. Inorg. Chem. 1991;
30: 4460–4464.
43. Nanthakumar A and Goff HM. J. Am. Chem. Soc.
1990; 112: 4047–4049.
44. Yu BS and Goff HM. J. Am. Chem. Soc. 1989; 111:
6558–6562.
45. Droghetti E, Nicoletti FP, Bonamore A, Sciamanna
N, Boffi A, Feis A and Smulevich G. J. Inorg. Bio-
chem. 2011; 105: 1338–1343.
1
973; 95: 6950–6956.
2. Boucher LJ. J. Amer. Chem. Soc. 1968; 90:
640–6645.
3. Watanabe J-Y and Setsune J-I. Organometallics
997; 16: 3679–3683.
6
1
4. Buchler JW and Rohbock K. Inorg. Nucl. Chem.
Lett. 1972; 8: 1073–1076.
5. Imamura T and Furusaki A. Bull. Chem. Soc. Jpn.
1
990; 63: 2726–2727.
6. Imamura T, Tanaka T and Fujimoto M. Inorg. Chem.
985; 24: 1038–1041.
1
7. Imamura T, Numatatsu T, Terui M and Fujimoto M.
Bull. Chem. Soc. Jpn. 1981; 54: 170–174.
8. Sishta C, Ke M, James BR and Dolphin D. J. Chem.
Soc., Chem. Commun. 1986: 787–788.
9. Buchler JW and Kruppa SB. Z. Naturforsch., B:
Chem. Sci. 1990; 45: 518–530.
0. Antipas A, Buchler JW, Gouterman M and Smith
PD. J. Am. Chem. Soc. 1980; 102: 198–207.
1. Buchler JW and Smith PD. Chem. Ber. 1976; 109:
46. Ghosh A and Taylor PR. J. Chem. Theory Comput.
2005; 1: 597–600.
47. Rots MJF and Zandstra PJ. Mol. Phys. 1982; 46:
1283–1310.
1
465–1476.
2. Liu W and Groves JT. Angew. Chem. Int. Ed. 2013;
2: 6024–6027.
48. Mun SK, Mallick MK, Mishra S, Chang JC and Das
TP. J. Am. Chem. Soc. 1981; 103: 5024–5031.
49. MeiningerDJ, CarantoJD,ArmanHDandTonzetich
ZJ. Inorg. Chem. 2013; 52: 12468–12476.
50. Fleischer EB and Srivastava TS. J. Am. Chem. Soc.
1969; 91: 2403–2405.
51. Adler AD, Longo FR, Kampas F and Kim J.
J. Inorg. Nucl. Chem. 1970; 32: 2443–2445.
52. Cheng RJ, Latos-Grazynski L and Balch AL. Inorg.
Chem. 1982; 21: 2412–2418.
53. Nanthakumar A and Goff HM. Inorg. Chem. 1989;
28: 4559–4561.
54. Bottomley LA and Kadish KM. Inorg. Chem. 1981;
20: 1348–1357.
55. Ellison MK, Schulz CE and Scheidt WR. Inorg.
Chem. 2000; 39: 5102–5110.
56. Crystal Clear; Rigaku/MSC Inc.: Rigaku Corpora-
tion, The Woodlands, TX, 2005.
5
3. Liu W, Huang X, Cheng M-J, Nielsen RJ, Goddard
WA and Groves JT. Science 2012; 337: 1322–1325.
4. Lee SC and Holm RH. Inorg. Chem. 1993; 32:
4
745–4753.
5. Jones JG, Tondreau GA, Edwards JO and Sweigart
DA. Inorg. Chem. 1985; 24: 296–300.
6. Scheidt WR, Lee YJ, Tamai S and Hatano K. J. Am.
Chem. Soc. 1983; 105: 778–782.
7. Gans P, Marchon JC and Moulis JM. Polyhedron
1
982; 1: 737–738.
8. Anzai K, Hatano K, LeeYJ and Scheidt WR. Inorg.
Chem. 1981; 20: 2337–2339.
9. Cohen IA, Summerville DA and Su SR. J. Am.
Chem. Soc. 1976; 98: 5813–5816.
0. Ogoshi H,Watanbe E,Yoshida Z, Kincaid J and Naka-
moto K. J. Amer. Chem. Soc. 1973; 95: 2845–2849.
1. Sadasivan N, Eberspaecher HI, Fuchsman WH and
Caughey WS. Biochemistry 1969; 8: 534–541.
2. Caughey WS and Johnson LF. J. Chem. Soc. D
57. ABSCOR; Higashi: Rigaku Corporation, Tokyo,
Japan, 1995.
58. Sheldrick GM. Acta Crystallogr., Sect. A 2008;
A64: 112–122.
59. SHELXTL97: Program for Refinement of Crystal
Structures; Sheldrick GM: University of Göttingen:
Göttingen, Germany, 1997.
1
969: 1362–1363.
3. Takahashi A, Yamaki D, Ikemura K, Kurahashi T,
Ogura T, Hada M and Fujii H. Inorg. Chem. 2012;
5
1: 7296–7305.
3
3
4. Kumar D, Sastry GN and de Visser SP. J. Phys.
Chem. B 2012; 116: 718–730.
5. TakahashiA, Kurahashi T and Fujii H. Inorg. Chem.
60. Laitar DS, Müller P, Gray TG and Sadighi JP.
Organometallics 2005; 24: 4503–4505.
61. Richmond TG. Coord. Chem. Rev. 1990; 105:
221–250.
2
011; 50: 6922–6928.
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 422–423

A CONVENIENT PROCEDURE FOR THE SYNTHESIS OF FLUORO-IRON(III) COMPLEXES
423
6
2. Sakai T, Ohgo Y, Hoshino A, Ikeue T, Saitoh T,
Takahashi M and Nakamura M. Inorg. Chem. 2004;
3: 5034–5043.
3. Moss TH, Bearden AJ and Caughey WS. J. Chem.
Phys. 1969; 51: 2624–2631.
4. Momenteau M, Mispelter J and Lexa D. Biochim.
Biophys. Acta, Gen. Subj. 1973; 320: 652–662.
5. Paeng IR, Shin JY, Nam H, Kim D-G and Paeng
K-J. Bull. Korean Chem. Soc. 1997; 18: 730–733.
6. Paeng IR, Nam H, Lee SW and Paeng K-J. J. Raman
Spectrosc. 1997; 28: 229–234.
69. Phillippi MA, Shimomura ET and Goff HM. Inorg.
Chem. 1981; 20: 1322–1325.
70. Ikezaki A, Takahashi M and Nakamura M. Dalton
Trans. 2011; 40: 9163–9168.
71. Hickman DL, Nanthakumar A and Goff HM. J. Am.
Chem. Soc. 1988; 110: 6384–6390.
72. Panchmatia PM, Ali ME, Sanyal B and Oppeneer
PM. J. Phys. Chem. A 2010; 114: 13381–13387.
73. Schulz CE, Song H, Mislankar A, Orosz RD, Reed
CA, Debrunner PG and Scheidt WR. Inorg. Chem.
1997; 36: 406–412.
74. Li M, Neal TJ, Wyllie GRA, Schulz CE and Scheidt
WR. Inorg. Chem. 2010; 49: 8078–8085.
75. Lang G, Boso B, Erler BS and Reed CA. J. Chem.
Phys. 1986; 84: 2998–3004.
4
6
6
6
6
6
6
7. Galich L, Huckstadt H and Homborg H. J. Porphyrins
Phthalocyanines 1997; 1: 259–265.
8. Kincaid J and Nakamoto K. Spectrosc. Lett. 1976;
9
: 19–21.
Copyright © 2014 World Scientific Publishing Company
J. Porphyrins Phthalocyanines 2014; 18: 423–423