Isolation and identification of the pre-catalyst in iron-catalyzed direct arylation of pyrrole with phenylboronic acid

Article abstract of DOI:10.1016/j.ica.2018.03.036

Herein we describe the synthesis, characterization, and role of three dichloric iron(III) complexes, [L1Fe(III)(Cl)₂]ClO₄ (L1Fe), [L2Fe(III)(Cl)₂]ClO₄ (L2Fe), and [L3Fe(III)(Cl)₂]ClO₄ (L3Fe

Full text of DOI:10.1016/j.ica.2018.03.036

Accepted Manuscript
Research paper
Isolation and Identification of the Pre-Catalyst in Iron-Catalyzed Direct Aryla-
tion of Pyrrole with Phenylboronic Acid
Samantha M. Brewer, Philip M. Palacios, Hannah M. Johnston, Brad S. Pierce,
Kayla N. Green
PII:
S0020-1693(18)30215-9
https://doi.org/10.1016/j.ica.2018.03.036
ICA 18181
DOI:
Reference:
Inorganica Chimica Acta
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Revised Date:
Accepted Date:
5 February 2018
21 March 2018
22 March 2018
Please cite this article as: S.M. Brewer, P.M. Palacios, H.M. Johnston, B.S. Pierce, K.N. Green, Isolation and
Identification of the Pre-Catalyst in Iron-Catalyzed Direct Arylation of Pyrrole with Phenylboronic Acid, Inorganica
Chimica Acta (2018), doi: https://doi.org/10.1016/j.ica.2018.03.036
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Isolation and Identification of the Pre-Catalyst in Iron-Catalyzed Direct
Arylation of Pyrrole with Phenylboronic Acid
Samantha M. Brewer,a Philip M. Palacios,b Hannah M. Johnston,a Brad S. Pierce,b and Kayla N. Greena*
a Department of Chemistry and Biochemistry, Texas Christian University, Fort Worth, TX, USA
b Department of Chemistry and Biochemistry, The University of Texas at Arlington, Arlington, TX,
USA
Abstract. Herein we describe the synthesis, charcterization, and role of three dichloric iron(III)
complexes, [L1Fe(III)(Cl)2]ClO4 (L1Fe), [L2Fe(III)(Cl)2]ClO4 (L2Fe), and [L3Fe(III)(Cl)2]ClO4 (L3Fe)
[L1 (Pyclen)=1,4,7,10-tetra-aza-2,6-pyridinophane; L2 =3,6,9,15-tetraazabicyclo[9.3.1]penta-deca-
1(15),11,13-trien-13-ol; L3 =3,6,9,15-tetra-azabicyclo[9.3.1]penta-deca-1(15),11,13-trien-12-ol], in the
coupling of pyrrole and phenylboronic acid to form 2-phenylpyrrole. The oxidation state and spin state
of the iron complexes were characterized using X-ray crystallography, UV-vis absorbance spectroscopy,
electron paramagnetic resonance spectroscopy, cyclic voltammetry, and mass spectrometry.
Electrochemistry results rank ligand L1-L3 as moderate tetra-azamacrocycle donors to iron between
cyclen and Me2EBC-12. Characterization of the iron(III) complexes and subsequent catalytic testing
indicates that the complexes enter the C-C coupling catalytic cycle in the high-spin iron(III) oxidation
state. Furthermore, the results indicate that the iron(III) complexes are essential for catalytic and
regioselective production of the 2-phenylpyrrole product.
KEYWORDS: Iron, cross-coupling, direct Suzuki-Miyaura, heterocycle, catalysis, pyrrole

Introduction
Carbon-Carbon cross-coupling reactions catalyzed by transition metals are invaluable components in a
chemist’s toolbox.[1] However, palladium is the most common metal used to facilitate these
transformations despite the low availability, toxicity, and cost of the precious metal.[1-7] Therefore, there
is a growing interest in replacing palladium with a more earth abundant element such as iron.[6-7] While
iron catalysts are indeed making strong contributions to the field of cross-coupling reactions,
mechanistic insights and thorough catalyst characterizations are much more challenging than the
palladium counterparts due to the strong reactivity of iron complexes, unpaired electrons complicating
NMR spectroscopy, and transient nature of intermediate species.[1-7] Nevertheless, iron complexes have
proven useful in the synthesis of organic compounds[1-18] and as model complexes of metalloenzymes.[19-
21]
Such work has resulted in reports that iron species, such as Fe(III)OOH, Fe(IV)=O, or Fe(V)=O play a
key role in oxidative iron catalysis and mechanistic details are becoming more understood.[17, 20, 22]
However, the studies that investigate iron catalysts for C-C bond formation focus largely on the scope of
the catalysts and less on the properties of the active metal center. A compliment to studies of substrate
scope would focus on the identification of catalyst oxidation state and spin state (which may be tuned
by the ligand scaffold[23-24]) needed for the desired organic transformation to take place. Thorough
understanding of ligand effects on the metal center will aid in designing more efficient catalysts.[1, 5-7]
For example, the White-Chen [Fe(S,S-PDP)] catalyst has been studied for C-H bond activation.[25] This
work focused on regioselectivity derived from substrate properties such as electronics, steric bulk, and
directing groups. Talsi et al. have separately probed the nature of the active species in this process using
EPR and enantioselectivity studies to show that the oxygen transfer occurs by an Fe(V)-oxo species.[26]
Insight into a previously unknown mechanism resulted from the detailed characterization in this

work.[27] Furthermore, multi-dentate N-containing ligands have been used to obtain stable Fe(IV)-oxo
[10, 19-20, 28-30]
species, providing invaluable spectroscopic comparisons to metalloenzymes in nature
.
Figure 1. Tetra-azamacrocycles studied by Wen et al. in combination with iron(II) salts to facilitate the
coupling of pyrrole and phenylboronic acid to produce 2-phenylpyrrole.[31]
Figure 2. Iron(III) complexes derived from cyclen, LN4H2, and Me2EBC-12.
Interesting, Bedford and co-workers reported that iron catalysts derived from rigid tetra-
azamacrocycles, such as Me2EBC-12, resulted in poor yields for the cross-coupling of 4-tolyl magnesium
bromide with cyclohexylbromide.[5] However in 2010, Wen. et al. reported that the tetra-
azamacrocycles shown in Figure 1, when mixed with iron(II) salts in the presence of oxygen, facilitate
direct arylation of pyrrole with phenylboronic acid to form 2-phenylpyrrole.[31] A preliminary mechanism
was proposed in which an iron-oxo species acts as the active catalyst; however, no metal oxidation
states were assigned or catalyst characterization reported aside from a mass spectrum that proved to be
[8, 11, 16, 24,
tenuous in its assignment. Since the release of this publication, it has been cited over 70 times
32-97]. Interestingly, to date the synthesis and characterization of iron complexes derived from two of the
four ligands in these original reports, LN4H2 and cyclen, have been reported (Figure 2). Both complexes
[(LN4H2)Fe(Cl)2]+ and [(cyclen)Fe(Cl)2]+ were identified as high-spin iron(III) systems.[98-101] Of the four

mixtures tested for catalytic ability by Wen and co-workers, the mixture containing L1 afforded the
highest yield. Therefore and reported here, we identified the spin-state and oxidation state of the
complex formed by L1 and iron(II) in the presence of oxygen and compared the structural and electronic
properties to [(LN4H2)Fe(Cl)2]+, [(cyclen)Fe(Cl)2]+, and others. We have previously explored L1 and its
derivatives (L2 and L3, Figure 3) as chelates for Cu(II), Ni(II), and Zn(II); the donor capacity of the ligand
was affected by the presence and position of the hydroxyl group.[102] Therefore, the iron complexes of
L2 and L3 were also isolated, characterized, and compared within the series described above. The bona
fide iron(III) high-spin complexes derived from L1, L2, and L3 were identified as pre-catalysts for direct
Suzuki-Miyaura coupling of pyrrole and phenylboronic acid to yield 2-phenylpyrrole. Finally, further
experiments show that some amine ligands can promote a small amount of background reactivity
yielding multiple products, but the iron(III) pre-catalysts are critical for focusing the reactivity to produce
only 2-phenylpyrrole, thus validating the need for the intact iron complex as a catalyst.
Figure 3. Chemical structure of ligands L1-L3 (
L1 = 1,4,7,10-tetra-aza-2,6-pyridinophane,[103-104]
L2 = 3,6,9,15-tetra-azabicyclo[9.3.1]penta-deca-1(15),11,13-trien-13-ol,[105] L3 = 3,6,9,15-tetra-
azabicyclo[9.3.1]penta-deca-1(15),11,13-trien-12-ol).[102]
Experimental Section
General Methods. Iron(II) perchlorate was freeze dried prior to use, all other reagents were
obtained from commercial sources and used as received, unless noted otherwise. NMR spectra
were obtained on a 400-MHz Bruker Advance spectrometer, using deuterated solvents (CDCl3).
NMR spectra were referenced using the corresponding solvent resonance (in parts per million;

CDCl3 δ = 7.26).[106] The following abbreviations were used for proper identification of the NMR
signals: s = singlet, d = doublet, t = triplet, m = multiplet. ESI-MS experiments were carried out
using an Agilent 6224 Accurate-Mass Time-of-Flight (TOF) mass spectrometer using 175 V to
ionize the complexes. Elemental analysis was performed by Canadian Microanalytical Service
Ltd. Electronic absorption spectra were recorded on a DU 800 UV-vis spectrophotometer
(Beckman Coulter) using a 3 mL quartz cuvette with a 1 cm path length. GC-MS analysis was
carried out using a Bruker Scion 436-GC-MS equipped with an auto sampler 8410 and a Br-5ms
column 29.9m. Caution! Perchlorate salts are explosive and should be handled in small
quantities. In particular, such compounds should never be heated as solids.
Synthesis of ligands. Stability of transition-metal complexes containing a tetra-azamacrocycle
are facilitated by the macrocyclic effect. The ease with which these processes occur depends
upon ring size, the number of donor atoms, electronic characteristics, and other factors.[104, 107]
The formation of 14-membered tetra-azamacrocyclic ligands, for example, typically involves the
use of transition metal-ions to template the cyclization step between two independent units to
form the ligand. The product of this reaction is, therefore, a transition metal macrocyclic
complex. However, tetra-azamacrocyclic amines, comprised of 12 atoms in the ring, form
through metal-independent cyclization pathways.[108-114] Therefore, for the work described
herein, the 12-membered pyridine and pyridol based tetra-azamacrocyclic amines ligands, L1-L3,
were produced previously reported procedures developed in our group and isolated as the
corresponding HCl salts.[102, 105] Caution! Perchlorate salts are explosive, and should be handled
in small quantities. In particular, such compounds should never be heated as solids.
[L1Fe(III)(Cl)2]ClO4 (L1Fe): Ligand (L1·3HCl) (101.1 mg, 0.3216 mmol) and Fe(ClO4)2 (83.7 mg,
0.330 mmol) were dissolved in 3 mL DI water; the solution was adjusted to pH=6 using 1M KOH.
The resulting red solution was allowed to stir open to air for 15 hours at 40oC. After 15 hours, a
tan precipitate was removed by centrifugation followed by filtration using a 0.45 m PTFE filter.
The water was removed using an azeotrope formed with acetonitrile. The resulting solid was
taken up in CH3CN and dried with Na2SO4. The addition of Et2O to the CH3CN solution, followed
by centrifugation yielded the product as a brown powder. Yield: 62% (92.8 mg, 0.198 mmol).
Yellow X-ray quality crystals were obtained by slow diffusion of ether into DMF at 4oC, CCDC#

1422489. ESI-MS (m/z) Found: 260.1515, [L1Fe(III)-2H+]+ (34%); 296.1360, [L1Fe(III)Cl-H+]+
,
,
(58%), 332.1209, [L1Fe(III)2Cl-]+, (14%). Theoretical: 260.0724, [L1Fe(III)-2H+]+, 296.0491,
[L1Fe(III)Cl-H+]+, 332.0258, [L1Fe(III)2Cl-]+. UV-vis, λmax, ε (M-1·cm-1): 261 nm (3,600), 311 nm
(800), 416 nm (170). Elemental analysis: [L1Fe(III)(Cl)2]ClO4 (Formula:C11H18N4FeO4Cl3); Found
(Calculated): C, 30.66 (30.50); H, 4.09 (4.20); N, 13.04 (12.95).
[L2Fe(III)(Cl)2]ClO4 (L2Fe): Ligand (L2·3HCl) (65.0mg, 0.208 mmol) was dissolved in 2.5 mL DI
water. The pH of the solution was adjusted to 5 using 1M KOH. Fe(ClO4)2 (53.6 mg, 0.211 mmol)
was dissolved in 1 mL DI water and added dropwise to the ligand solution; the pH was
maintained between 3.5 and 5.2. After all iron(II) solution was added, the pH was adjusted to
5.3. The solution was allowed to stir open to air 2 days resulting in precipitation of a brown solid
that was isolated by centrifugation. Yield: 30% (27.4 mg, 0.061 mmol). Yellow crystals suitable
for XRD analysis were obtained by slow evaporation of water at room temperature, CCDC #
1422490. ESI-MS (m/z) Found: 276.1529, [L2Fe(III)-2H+]+ (35%); 312.1377, [L2Fe(III)Cl-H+]+ (45%).
Theoretical: 276.0674, [L2Fe(III)-2H+]+ 312.0440 [L2Fe(III)Cl-H+]+. UV-vis, λmax, ε (M-1·cm-1): 249
;
nm (6,700), 306 nm (4,000), 356 nm (2,700). Elemental analysis: [L2Fe(III)(Cl)2]ClO4 (Formula:
C11H18N4FeO5Cl3); Found (Calculated): C, 29.88 (29.46); H, 4.08 (4.05); N, 11.75 (12.49).
[L3Fe(III)(Cl)2]ClO4 (L3Fe): Ligand (L3·3HCl) (206.4 mg, 0.6223 mmol) was dissolved in 20 mL DI
water and the pH was adjusted to 5 using 1M KOH. Fe(ClO4)2 (164.0 mg, 0.6155 mmol) was
dissolved in 10 mL DI water, added drop wise to ligand, and the pH was re-adjusted to 5. The
solution was allowed to stir open to air for 2 days at 40oC. The solvent was removed under
reduced pressure. The resulting dark brown solid was dissolved in DMF, dried with Na2SO4, and
filtered. Ether was added to the DMF solution and a red-brown powder was isolated by
centrifugation. Yield = 63.2% (235.6 mg, 0.3932 mmol). Brown crystals suitable for XRD analysis
were obtained by slow evaporation from water at room temperature, CCDC # 950048. ESI-MS
(m/z) Found: 276.1481, [L3Fe(III)-2H+]+ (30%), 312.1322, [L3Fe(III)Cl-H+]+ (55%). Theoretical:
276.0674, [L3Fe(III)-2H+]+, 312.0440 [L3Fe(III)Cl-H+]+. UV-vis, λmax, ε (M-1·cm-1): 205 nm (14,000),
219 nm (10,000), 284 nm (5,600), 458 nm (300). Elemental analysis: [L3Fe(III)(Cl)2]ClO4 (Formula:
C11H18N4FeO5Cl3); Found (Calculated): C, 29.03 (29.46); H, 4.05 (4.05); N, 12.57 (12.49).

X-ray Diffraction Analysis. Crystal diffraction data were collected at 100 K on a Bruker D8 Quest
Diffractometer. Data collection, frame integration, data reduction (multi-scan), and structure
determination were carried out using APEX2 software.[115] Structural refinements were
performed with XSHELL (v 6.3.1), by the full-matrix least-squares method.[116] All non-hydrogen
atoms were refined using anisotropic thermal parameters, while the hydrogen atoms were
treated as mixed. The ORTEP molecular plots (50 %) were produced using APEX2 (Version
2014.9-0).
Electrochemistry. Cyclic voltammetry experiments were obtained using 2.2 mM complex and
100 mM tetrabutylammonium tetraflouroborate as the supporting electrolyte in DMF. The
electrochemical cell was composed of a working glassy carbon electrode, a Pt auxiliary
electrode, and a silver wire as the reference electrode. To facilitate solubility, all samples were
first dissolved in 1M HCl, thoroughly dried, and then dissolved in DMF for electrochemical
analysis. The potential was scanned in the negative direction at a rate of 100 mV/s, starting at
the open circuit potential. The potential values presented here have been normalized to the
half-wave potential of the Fc/Fc+ redox couple set equal to 0.00 V. For comparison purposes
half-wave potential in cited references were converted to reflect Fc/Fc+ = 0 mV; [(Me2EBC-
12)Fe(Cl)2]PF6, (Fc/Fc+ = 400 mV), and [(cyclen)Fe(Cl)2]Cl, (Fc/Fc+ = 515 mV).[99, 117-118]
X-band EPR Spectroscopy and Analytical Simulations: X-band (9 GHz) EPR spectra were
recorded on a Bruker EMX Plus spectrometer equipped with a bimodal resonator (Bruker model
4116DM). Low-temperature measurements were made using an Oxford ESR900 cryostat and an
Oxford ITC 503 temperature controller. A modulation frequency of 100 kHz was used for all EPR
spectra. All experimental data used for spin-quantitation were collected under non-saturating
conditions. Analysis of the EPR spectra utilized the general spin Hamiltonian,
Equation 1
where D and E are the axial and rhombic zero-field splitting parameters and g is the g-tensor.[119]
EPR spectra were simulated and quantified using Spin Count (ver. 5.8.6218.29549), written by
Professor M. P. Hendrich at Carnegie Mellon University. The simulations were generated with

consideration of all intensity factors, both theoretical and experimental, to allow concentration
determination of species. The only unknown factor relating the spin concentration to signal
intensity was an instrumental factor that depended on the microwave detection system.
However, this was determined by the spin standard, Cu(EDTA), prepared from a copper atomic
absorption standard solution purchased from Sigma–Aldrich.
2-phenylpyrrole Yield Determination: Phenylboronic acid (24 mg, 0.2 mmol) and crystalline
material of the iron complex (0.02 mmol) were added to a 5 or 10 mL flask equipped with a stir
bar. Pyrrole (1 mL) was added to flask, the mixture was heated to 130oC for 10 hours. The
reaction was cooled to room temperature and the pyrrole was removed under vacuum until no
visible liquid was present. Increasing the time the reaction was kept under reduced pressure
decreased yields. The product mixture was dissolved in a minimum amount of CDCl3, 5 µL of
dimethyldiphenylsilane was added to the solution. The solution was filtered through a 0.2 µm
nylon filter and a known amount of sample was added to a pre-weighed NMR tube. Yield
determinations were performed using three resonances 6.875, 6.532, and 6.307 ppm
corresponding to 2-phenylpyrrole and a resonance at 0.533 ppm corresponding to
dimethyldiphenylsilane. The reported values are averages of all resonances; each measurement
was run in triplicate.
Control Reactions: Phenylboronic acid (24 mg, 0.2 mmol) and ligand (0.02 mmol), if used, were
added to a 2 mL flask equipped with a stir bar, the system was then placed under an atmosphere
o
of nitrogen. Pyrrole (1 mL) was added to flask and the mixture was heated to 130 C for 15
minutes, if used, 10 mL O2 was injected directly into the pyrrole, the system was closed, and
heated for 10 hours. Yields were determined as stated above.
GC-MS Details: Method, 80C 2 min, ramp 5C/min to 170C, ramp 20C/min to 300C, hold 5
min. GC-MS Compound Identification L3 coupling: 3-methyl-4-phenylfuran-2(5H)-one: (trace) RT.
14.909 min. Found (Cal.): M+· 174.1 (174.1); (6%) 3-phenylpyrrole: RT. 15.002 min. Found (Cal.)
M+· 142.9 (143.0); (trace) 3-methyl-5-phenylfuran-2(4H)-one: 16.330 min. Found (Cal.) 174.1
(174.1); (16%) 2-phenylpyrrole: RT: 16.460 min Found (Cal.): M+· 142.9 (143.0).

Results and Discussion
Synthesis and Characterization
The corresponding iron complexes of L1-L3, shown in Figure 4, were synthesized in water at pH
~5 to compensate for the protonation of the isolated ligands. The iron(II) salt was used in
metalation of L1 L3 and was oxidized in air to iron(III) prior to isolation of the L1Fe L3Fe
-
-
complexes. The iron(II) perchlorate salt was exploited to facilitate the growth of X-ray quality
crystals, discussed below. Attempts to form complex using Fe(ClO4)3 did not afford product in
water. The presence and position of the hydroxyl group on the aromatic ring of the ligand
affected the metalation efficiency and solubility of the resulting complexes; therefore, divergent
synthetic strategies were developed for each complex produced. For example, mixtures of
CH3CN/Et2O or DMF/Et2O were used to isolate metal complexes L1Fe (62% yield) and L3Fe (63%
yield), respectively, as solids precipitates. L2Fe was easily isolated as a precipitate from the
aqueous reaction mixture, albeit with a low yield (30%). Nevertheless, the resulting ferric
complexes of L1-L3 were stable to both air and light and can be stored indefinitely once isolated
as dark red (L3Fe) or light brown (L1Fe, L2Fe) solids.
Figure 4. Pictorial representation of [L1Fe(III)(Cl)2]+
[L3Fe(III)(Cl)2]+ (L3Fe).
( (L2Fe), and
L1Fe), [L2Fe(III)(Cl)2]+
X-ray Crystallography
Figure 5 shows the results of single crystal diffraction analysis on crystalline solids of L1Fe, L2Fe,
and L3Fe. Yellow, X-ray quality crystals of L1Fe were obtained through vapor diffusion of ether
into a solution of DMF; yellow L2Fe and brown L3Fe crystalline materials were isolated by slow
evaporation of aqueous solutions. Table S1 contains the crystal data, intensity collections, and
structure refinement parameters; a full list of bond lengths and angles are also located in the
supporting information. The structures determined through X-ray diffraction analysis show that

each complex adopts a six coordinate, distorted octahedral geometry (N(2)-Fe-N(4), ~85o; N(1)-
Fe-N(3), ~147o). The coordination sphere consists of four nitrogen donors from the ligand set
and two chloride ions. Each complex adopts a cis-folded geometry due to the rigidity of the 12-
membered ligand set in which one chloride is cis and the other is trans to the pyridol ring.16 This
finding is consistent with previous structure determination of [L1Fe]BF4, reported by Alcock et
al.[120] The charge of the [LxFe(III)(Cl)2]+ systems are balanced by one perchlorate counter ion
within the unit cell. The position of the hydroxyl group and the flexibility of the aliphatic portion
of the ligand results in an enantiomeric mixture of L3Fe. The hydroxyl group was modeled for
disorder to account for both enantiomers, as shown in Figure 5c.
Figure 5. ORTEP (50%) representations of L1Fe (A), L2Fe (B), and L3Fe (C). The perchlorate anion has
been omitted for clarity; modelling of disorder for L3Fe is shown in grey. All complexes take on a cis-
folded distorted octahedral geometry. The Fe-N bond lengths are greater than 2.0 Å, consistent with
other high-spin ferric systems.[99, 101, 118] A full list of bond lengths and angles are available in Tables S2-
7.

Table 1. Selected bond lengths (Å) and angles (o) of complexes L1Fe
-
L3Fe
.
Bond
L1Fe
L2Fe
L3Fe
Fe(1)-N(1)
Fe(1)-N(2)
Fe(1)-N(3)
Fe(1)-N(4)
2.1641(6) 2.1787(11) 2.162(2)
2.2001(6) 2.2023(10) 2.181(2)
2.1676(6) 2.1812(11) 2.172(2)
2.1074(5) 2.0970(10) 2.107(2)
N(1)-Fe(1)-N(3) 147.25(2) 146.70(4) 147.17(8)
N(2)-Fe(1)-N(4) 85.56(2) 85.47(4) 85.72(8)
The geometry of complexes L1Fe-L3Fe can be compared to other macrocycles in the literature.
For example, N(1)-Fe-N(3) bond angles have been reported for two 12-membered macrocycles
(Figure 2), [(LN4H2)Fe(Cl)2]+ (142.41°) and [(cyclen)Fe(Cl)2]+ (146.40°), where the latter complex
provides the closest bond angle to the L1Fe-
L3Fe systems.[101, 117, 121] The L2Fe complex provides
slightly longer Fe-N bonds compared to L1Fe and L3Fe. For example, the Fe-N(2) bond of L2Fe
was determined to be 2.2023(10) Å, while L1Fe and L3Fe were slightly shorter with 2.2001(6)
and 2.181(2) Å, respectively. The only exception was observed with the Fe-N(4) bond, which was
slightly shorter in L2Fe (2.0967(10) Å) compared to L1Fe (2.1074(5) Å) and L3Fe (2.107(2) Å). The
difference in the equatorial Fe-N(4) pyridine-derived bond length is consistent with a stronger
interaction between the pyridol nitrogen of L2 vs. L1 and L3. Throughout the series, the
equatorial Fe-N(2) (pyridine atom) bond is the longest and thus the weakest Fe-N interaction.
The Fe-N bond lengths of L1Fe-L3Fe are greater than 2.00 Å (Tables 1-2 and Figure 5) and are
consistent with an iron(III) high-spin system.[122] For example, the iron center of [(LN4H2)Fe(Cl)2]+
was assigned as a high-spin iron(III) by EPR, and the Fe-N bond lengths within the high-spin
complex were measured as 2.128 and 2.221 Å.[117, 123] Altogether, the results indicate that
ligands L1
donor to iron(III) compared to L1 and L3. However, the differences in bond lengths and angles
within L1Fe L3Fe is much smaller than the nickel(II), copper(II), and zinc(II) congeners of L1
L3 [102]
-L3 stabilize the high-spin iron(III) in similar manners and that L2 is a slightly stronger
-
-
.

Mass Spectrometry
Mass spectrometry further confirmed the oxidation state of the iron(III) in L1Fe-L3Fe. The mass
spectrum obtained for L1Fe consisted of three isotopic envelopes that correspond to the
complex: 260.1519 m/z = [L1Fe(III)-2H+]+, 296.1358 m/z = [L1Fe(III)Cl--H+]+, and 322.1209 m/z =
[L1Fe(III)2Cl-]+ (Figure S4). Similar fragmentation patterns were obtained for complexes L2Fe and
L3Fe and are detailed in experimental methods related to each complex. Mass spectrometry
analysis of the coupling reaction performed by Wen. et al. revealed three isotopic envelopes: m/z =
207.1532, 353.1956, and 369.1956.[31] The isotopic envelopes were assigned as [L1]+, [L1 + Fe +C2O4]+,
and [L1 + Fe + C2O4 + O]+, respectively, with no indication of iron oxidation states or charge balance. The
isotopic envelope observed at m/z = 207.1532 indeed corresponds to the singly protonated free ligand
[L1 + H+]+, which we observe as well with studies of free ligand L1. However, the assignment of m/z =
353.1965 as [L1 + Fe +C2O4]+ is incorrect, as the exact mass the expected species is modeled to have m/z
= 350.0672, three mass units less than the observed ion reported. Similarly, the assignment of m/z =
369.1897 as [L1 + Fe + C2O4 + O]+ (Theoretical m/z = 366.0621) as a component of the catalytic reaction
does not correlate as well. Therefore, the results reported herein serve as the first validation of the
composition and oxidation state of the iron pre-catalyst involved in the C-C coupling chemistry, to be
described later.
Spin State Determination
The spin and oxidation states of complexes L1Fe
via electron paramagnetic resonance (EPR) spectroscopy. As shown in Figure 6 (left), the EPR
spectra (solid lines) for all complexes (L1Fe L3Fe) exhibit features typical of high-spin ferric iron
-L3Fe were also validated at low temperature
-
(S = 5/2). For analytical purposes, all data were recorded under non-saturating microwave
power. The simulations overlaid on each spectrum (dashed lines) consist of contributions from
two separate doublets. As indicated by the energy diagram shown in Figure 6 (right), the
dominant transition for these complexes arises from the ground ms = ± 1/2 doublet of a S = 5/2
spin state with near axial symmetry (E/D = 0.07). Transitions within this doublet yield the
observed g-values of 7.6, 4.3, and 1.7. The linewidth of this transition can be reasonably
simulated by assuming a Gaussian distribution in rhombicity (E/D) [designated σE/D], which
broadens the g ~ 1.7 resonance significantly. The lower intensity features observed at g ~ 5.8
and 1.97 are nearly absent at low temperature (4 K) but reach a maximal intensity near ~8 K

before decreasing again as temperature approaches 20 K. The alternating temperature
dependence of these features confirm that this signal must originate from the middle ms = ± 3/2
doublet of the S = 5/2 spin state. The magnitude of the zero-field splitting parameter (|D| = 0.7 ±
0.2 cm-1) was determined by plotting the EPR signal intensity of the ms = ± 1/2 doublet versus
1/T and fitting the data to a Boltzmann population distribution for a 3-level system. Additional
corroboration of the axial zero-field splitting term was obtained by simultaneous simulation of
EPR spectra collected at temperatures ranging from 4 to 20 K (n = 5). Within this temperature
regime, all simulations accurately reproduce the relative intensity for each transition (± 1/2 and
± 3/2 ms-states) using a D-value of 0.7 ± 0.2 cm-1. Within error, all complexes (L1Fe
- L3Fe)
exhibited equivalent temperature dependence and thus all EPR simulations shown in Figure 6
utilized the same axial zero field splitting term. Indeed, with the exception of minor
perturbations in the extent of E/D-distribution (σE/D), all complexes L1Fe
– L3Fe exhibit nearly
equivalent EPR spectroscopic properties. The near equivalent of EPR spectra observed for these
complexes is understandable given the close agreement in Fe-coordination sphere bond length
and coordination geometry observed crystallographically (Table 1).
Table 2. Comparison of spin-state and bond lengths within iron(III) complexes derived from 12-
membered tetra-azamacrocyclic ligands.11,12
Complex
Spin State (S) Fe-N4 (Å)
Fe-N2 (Å)
Ref.
N1-Fe-N3 ()
[99]
[(cyclen)Fe(Cl)2]Cl
[(LN4H2)Fe(Cl)2]Cl
5/2
-
2.1461(16) 146.40(6)
[101, 117, 121]
[118]
2.094(1)
-
2.189(1)
2.163 (2)
2.2001(6)
142.41(7)
153.20

[(Me2EBC-12)Fe(Cl)2]PF6 5/2
[(L1)Fe(Cl)2]ClO4
[(L2)Fe(Cl)2]ClO4
[(L3)Fe(Cl)2]ClO4
5/2
5/2
5/2
2.1074(5)
147.25(2)
‡
‡
‡
2.0967(10) 2.2023(10) 146.71(4)
2.107(2) 2.181(2) 147.19(8)
* not reported, ‡ This work

Figure 6. X-band EPR spectra of L1Fe
- L3Fe (left). Quantitative simulations (dashed lines) are
overlaid on each spectrum for comparison. The black circle observed at g ~ 4.92 in L2Fe is from a
minor (< 10%) high-spin iron(III)-impurity. Instrumental parameters: frequency, 9.643 GHz,
microwave power, 6 μW; modulation amplitude, 0.9 mT; temperature, 10 K. Simulation
parameters: S = 5/2; g1,2,3 ~2.0; |D|, 0.7 ± 0.2 cm-1; E/D, 0.07; σE/D, 0.01; σB, 0.9 mT. Energy level
diagram (right) illustrating the splitting of doublets within the S = 5/2 spin state along each
principle axis (X, red; Y, green; Z, black).
Electrochemistry
Cyclic voltammetry was used to evaluate the electrochemical behavior of L1Fe
-L3Fe. The cyclic
voltammograms corresponding to the iron(III/II) couple of L1Fe L3Fe are shown in Figure 7. Of
-
the three complexes, L2Fe (E1/2 = -486 mV) has the most negative half potential, followed by
L3Fe (E1/2 = -468 mV) and L1Fe (E1/2 = -465 mV). The difference in the half potentials indicated
derivatization of the pyridine ring affects the electron density around the iron center,
specifically, L2 is the most donating. Figure 8 compares the electrochemical potentials of iron
complexes containing 12-membered tetra-azamacrocycles; a wide range (-865 mV to 20 mV) of
half potentials is achieved by changing the donor capacity of the ligand set, LN4Me2
L1 L3 L2
cyclen 94, 105, 122,[124]
< Me2EBC-12
<
≈
<
<
.
The reversibility of the redox process was investigated by determining the ΔEp and Ipa/Ipc. The
iron(III/II) redox processes are quasi-reversible (ΔEp, Ipa/Ipc): L1Fe (105 mV, 1.2583), L2Fe (112
mV, 0.7264), and L3Fe (101 mV, 0.7485). Lastly, the electrochemical events are diffusion

controlled for all three iron complexes, as shown by the linear relationship between Ip and the
square-root of the scan rate (Figure S6).[125]
It should be noted that an additional ligand based oxidation event around 900 mV is observed in
the full solvent window (1.4 to -1.2 mV, Figure S5) for L1Fe
that this event was ligand derived, based on electrochemical analysis of the corresponding
zinc(II) complexes providing a similar behavior.[102] Electrochemical analysis of L1
L3 in DMF
-L3Fe. We have previously postulated
-
solvent with TBAP electrolyte provide direct confirmation that this positive oxidation wave is
ligand based (Figure S5).
10 µA
L1Fe
L2Fe
L3Fe
300
-200
-700
-1200
Potential (mV) vs. Fc/Fc+
Figure 7. Cyclic voltammogram overlay of the FeIII/II couple measured for L1Fe
-L3Fe in DMF
containing 0.1 M [Bu4N][BF4] as electrolyte, Ag/Ag+ reference electrode, glassy carbon working
electrode, and platinum auxiliary electrode at a scan rate of 100 mV/sec. All scans were
referenced to Fc/Fc+ = 0.00 mV.
Figure 8. Iron(III)/(II) halfway potentials of iron complexes in literature containing 12-
membered tetra-azamacrocycles. The potentials are reported as referenced to Fc/Fc+ =
0.00 mV. [99, 117-118, 124, 126]

Table 3. The anodic wave potential (Epa), cathodic wave potential (Epc), peak potential
separation (ΔEp), and halfway potential (E1/2) of L1Fe-L3Fe. L2Fe contains the most stable
iron(III) ion in the series according to the E1/2 values.
Ipa/Ipc
Complex Epc (mV) Epa (mV) E1/2 (mV) ΔEp (mV) Ipc (A) Ipa (A)
L1Fe
L2Fe
L3Fe
-517
-542
-519
-412
-430
-418
-465
-486
-468
105
112
101
10.5837 -13.3181 1.2583
18.4176 -13.3791 0.7264
11.9661 -8.9570 0.7485
Figure 9. The electronic absorbance spectra of ligands L1-L3 and complexes L1Fe-
L3Fe obtained in 1 M HCl.
Electronic Absorption Spectroscopy
The spectrophotometric behaviors of L1Fe-L3Fe are shown in Figure 9. The aromatic
components of the ligands (sans metal) result in absorbance bands between 210 and
290 nm when measured in 1 M HCl. The iron(III) complexes showed no appreciable
shift in the ππ* region (210-300 nm). Instead, the appearance of MLCT bands were
observed at wavelengths greater than 300 nm: L1Fe (311 nm), L2Fe (306 and 356
nm), L3Fe (458 nm), which is consistent with high-spin d5 complexes.[127] These metal
based assignments are supported by comparison to the 12-membered macrocycle,
[(cyclen)Fe(NCMe)2]3+ that Hua et al. reported to have absorbance bands at 259 and
358 nm and corresponding extinction coefficients below 500 M-1cm-1 [127]
.

Furthermore, the difference in the absorbance spectra of L1Fe, L2Fe, and L3Fe is
reflected in the visible color of the complexes. The L1Fe and L2Fe complexes are light
brown solids, while L3Fe is a red solid; in solution L1Fe and L2Fe are yellow and L3Fe
is brown. Low solubility of the complexes in other solvents precluded a full study of
solvent effects or at other pH values. Therefore, it should be noted that the iron
complex responsible for catalytic activity may vary in the degree of ligand
protonation, but the coordination sphere around the iron within the complex does
not change. The electronic absorption differences between L3Fe compared to the
L1Fe and L2Fe complexes could be attributed to the lack of symmetry originating
from the hydroxyl moiety in the meta-position of the pyridine ring of the ligand.
Table 4. Catalytic efficiency of iron complexes to obtain 2-phenylpyrrole from pyrrole
and phenylboronic acid, achieved using 10% catalyst loading in the presence of air.
Catalyst Yield
L1Fe
L2Fe
L3Fe
57 ± 3
58 ± 7
52 ± 7
Catalytic Activity of Iron(III) Complexes
Motivated by the report that addition of a tetra-azamacrocycle, iron(II) salt, and
oxygen to phenylboronic and pyrrole results in the formation of 2-phenylpyrrole[31],
we explored the oxidation state of the pre-catalyst by testing L1Fe, L2Fe, and L3Fe for
catalytic activity. Yields are shown in Table 4. In this series of experiments each
catalyst was tested at 10% loading, open to air. The ferric complexes afforded 2-
phenylpyrrole in yields of 57% (L1Fe), 58% (L2Fe), and 52% (L3Fe), thereby identifying
the oxidation state of the pre-catalyst as an iron(III) species. The realization that the
iron(III) complexes enter the catalytic cycle will allow for better foundation to
determine the oxidation state and identity of the active catalytic species. The
following discussion focuses on the experiments used to validate that the iron +

ligand catalyst species is solely responsible for providing the selective reactivity
observed, showing the ligand can catalyze baseline reactions with no regioselectivity,
and components necessary for the reaction to proceed.
Table 5. Control reactions used to determine the yield of product in the absence of
the high-spin iron(III) complexes.
Test Compound Oxidant
Yield
*
*
0
0
0
0
*
20 eq. oxygen
20 eq. oxygen
20 eq. oxygen
Fe(ClO4)3
L1
L2
20 eq. oxygen Trace
20 eq. oxygen Trace
Atmosphere Trace
L3
L3
*Not present
Control reactions were performed to ensure that the catalytic reactivity observed
with L1Fe
,
L2Fe, and L3Fe was due only to the iron complexes. The substrates,
o
pyrrole and phenylboronic acid, were heated to 130 C in both the absence or
presence of oxygen (Table 5). No reaction was observed under either of these
conditions. This indicates that the reaction requires a catalyst to proceed.
Additionally, four control reactions consisting of iron(III) perchlorate, L1, L2, and L3
were performed in the presence of 10 mL O2. Yield of 2-phenylpyrrole was
determined by both GC-MS and NMR due to the low quantities observed. Iron(III)
perchlorate and L1 did not afford 2-phenylpyrrole, however, L2 and L3 produced
trace amounts of 2-phenylpyrrole. Interestingly GC-MS analysis of the control
reaction, which included pyrrole, phenylboronic acid, L3, and atmospheric air also
showed the formation of trace amounts of 2-phenylpyrrole, 3-phenylpyrrole, and two
butenolides. The formation of 3-phenylpyrrole and butenolides was previously
observed by Campi et al. when 3-phenylprop-2-yn-1-amine is exposed CO/H2 at 400
psi at 70oC for 20 hours in the presence of a rhodium catalyst. These results indicate

that the ligand is capable of background reactivity thus producing pyrrole derivative.
Additionally, it indicates that the derivatives produced are controlled by the
composition of the oxidant (O2 and CO2) used in the reaction. Importantly,
comparison of the products formed in the presence of L3 and L3Fe and atmosphere
show that the use of the iron complex is essential to obtain catalytic and
regioselective production of 2-phenylpyrrole product.
Conclusion
The addition of iron(II) perchlorate to the tetra-azamacrocycles L1, L2, and L3 in the
presence of oxygen yields high-spin iron(III) complexes L1Fe, L2Fe, and L3Fe as shown
by X-ray crystallography and EPR spectroscopy. Furthermore, it was demonstrated
that the high-spin iron(III) complexes participate in the coupling of pyrrole and
phenylboronic acid to produce 2-phenylpyrrole. Although a small amount of
background reactivity was overserved with L3, the results show that the iron
complexes are responsible for controlling the reaction to produce 2-phenylpyrrole
alone. A full study involving a large library of iron macrocyclic derived complexes
focused on understanding features of catalytic activity is the topic of a forth coming
report. Finally, the characterization of the complexes using electrochemistry, UV-vis
spectroscopy and mass spectrometry lays a foundation for mechanistic investigations
concerning the oxidation state of the iron center throughout the catalytic process.
Supporting Information
Crystallographic data for L1Fe, L2Fe, and L3Fe (CCDC #: 1422489, 1422490, and
950048, respectively) in both CIF and table format.
Conflicts of interest
There are no conflicts to declare.
Acknowledgements
The authors are grateful for generous financial support from TCU Andrews Institute
of Mathematics & Science Education, TCU Research and Creativity Activity Grant, TCU

Invests in Scholarship, and INFOR Moncrief Foundation Support. The authors would
like to acknowledge Kimberly M. Lincoln for preliminary work on the iron complexes.
KG thanks the Online Research Discussions of Methods in Organometallic Catalysis
(ORDMOC) for helpful feedback related to this project.
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Highlights
 Synthesis of three new dichloric iron(III) tetra-aza macrocyclic complexes.
 Characterized via X-ray crystallography, UV-vis absorbance spectroscopy, electron
paramagnetic resonance spectroscopy, cyclic voltammetry, and mass spectrometry.
 Pre-catalyst for the coupling of pyrrole and phenylboronic acid to form 2-
phenylpyrrole.
 Iron(III) state of the complexes is essential for catalytic and regioselective production
of the 2-phenylpyrrole product.