A combined experimental and theoretical study of dinitrosyl iron complexes containing chelating bis(diphenyl)phosphinoX (X = benzene, propane and ethylene): X-ray crystal structures and properties influenced by the presence or absence of π-bonds in chelating ligands

Article abstract of DOI:10.1016/j.poly.2013.12.019

Recent discoveries involving the roles of nitric oxide in humans have stimulated intense interest in transition metal nitrosyl complexes. A series of dinitrosyl iron complexes with the formula [(DPPX)Fe(NO)₂], {DPPX = 1,2-bis(diphenylphosphino)benzene (1), 1,3-bis(diphenylphosphino)propane (2), and cis-1,2-bis(diphenylphosphino)ethylene (3)} has been prepared and characterized through a combination of FT-IR, NMR, UV-Vis, X-ray crystallography, and electrochemical techniques. Infrared spectroscopy showed NO shifts to the region of 1723 and 1674 cm^-1 for complexes 1 and 3, and 1708 and 1660 cm^-1 for 2, indicating that ligand 2 acts as a stronger σ-donor. The X-ray crystallographic data showed that 1 and 3 possess the rare repulso conformation while 2 has the attracto conformation. CV studies on compounds 1, 2 and 3 display two quasi-reversible oxidations with the E _1/2 values at 0.101 and 0.186 V, 0.121 and 0.184 V, and 0.019 and 0.342 V, respectively. The larger ΔE value for compound 2 compared with that of 1 and 3 is attributed to the lack of π-bonds between the two phosphorus atoms. Theoretical calculations using density functional theory were carried out on the synthesized compounds and model compounds and the results are consistent with the experimental data. The calculated HOMO-LUMO gaps for compounds 1, 2 and 3 are 3.736, 4.060, and 3.669 eV, respectively, which supports the stronger back-donation for compound 2 than that of compounds 1 and 3.

Full text of DOI:10.1016/j.poly.2013.12.019

Accepted Manuscript
A combined experimental and theoretical study of dinitrosyl iron complexes
containing chelating bis(diphenyl)phosphinoX (X = benzene, propane and eth-
ylene): X-ray crystal structures and properties influenced by the presence or
absence of π-bonds in chelating ligands
Lauren R. Holloway, Andrew J. Clough, Jessica Y. Li, Emily L. Tao, Fu-Ming
Tao, Lijuan Li
PII:
S0277-5387(13)00829-2
DOI:
http://dx.doi.org/10.1016/j.poly.2013.12.019
POLY 10473
Reference:
Polyhedron
To appear in:
Received Date:
Accepted Date:
22 October 2013
19 December 2013
Please cite this article as: L.R. Holloway, A.J. Clough, J.Y. Li, E.L. Tao, F-M. Tao, L. Li, A combined experimental
and theoretical study of dinitrosyl iron complexes containing chelating bis(diphenyl)phosphinoX (X = benzene,
propane and ethylene): X-ray crystal structures and properties influenced by the presence or absence of π-bonds in
chelating ligands, Polyhedron (2013), doi: http://dx.doi.org/10.1016/j.poly.2013.12.019
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A combined experimental and theoretical study of dinitrosyl iron complexes
containing chelating bis(diphenyl)phosphinoX (X = benzene, propane and ethylene):
X-ray crystal structures and properties influenced by the presence or absence of π-
bonds in chelating ligands
Lauren R. Hollowaya, Andrew J. Clougha, Jessica Y. Lib,
Emily L. Tao b, Fu-Ming Taob*, and Lijuan Lia*
a Department of Chemistry and Biochemistry, California State University, Long Beach,
1250 Bellflower Blvd., Long Beach, CA 90840 USA
b Department of Chemistry and Biochemistry, California State University, Fullerton, P.O.
Box 6866, Fullerton, CA 92834 USA
Keywords: dinitrosyl iron complexes, chelating bisphosphines, spectroscopy, DFT
calculations
*To whom correspondence should be addressed. aL.L.: tel, (562) 985-5068; email,
lli@csulb.edu. bF.M.T.: tel, (657) 278-4517; email, ftao@exchange.fullerton.edu.
1

Abstract:
Recent discoveries involving the roles of nitric oxide in humans have stimulated
intense interest in transition metal nitrosyl complexes. A series of dinitrosyl iron
complexes with the formula [(DPPX)Fe(NO)2], {DPPX = 1,2-
bis(diphenylphosphino)benzene (1), 1,3-bis(diphenylphosphino)propane (2), and cis-1,2-
bis(diphenylphosphino)ethylene (3)} has been prepared and characterized through a
combination of FT-IR, NMR, UV-vis, X-ray crystallography, and electrochemical
techniques. Infrared spectroscopy showed NO shifts to the region of 1723 and 1674 cm-1
for complexes 1 and 3, and 1708 and 1660 cm-1 for 2, indicating that ligand 2 acts as a
stronger σ–donor. The X-ray crystallographic data showed that 1 and 3 possess the rare
repulso conformation while 2 has the attracto conformation. CV studies on compounds
1, 2 and 3 display two quasi-reversible oxidations with the E°1/2 values at 0.101 and 0.186
V, 0.121 and 0.184 V, and 0.019 and 0.342 V, respectively. The larger ΔE value for
compound 2 compared with that of 1 and 3 is attributed to the lack of π-bonds between
the two phosphorus atoms. Theoretical calculations using density functional theory were
carried out on the synthesized compounds and model compounds and the results are
consistent with the experimental data. The calculated HOMO-LUMO gaps for
compounds 1, 2 and 3 are 3.736, 4.060, and 3.669 eV, respectively, which supports the
stronger back-donation for compound 2 than that of compounds 1 and 3.
2

Introduction:
The realization that nitric oxide (NO) is a biological messenger in many
physiological processes has brought about a renewed interest in its chemistry, particularly
the iron complexes that are central to the role of nitric oxide within the body.i Specifically,
non-heme iron nitrosyl complexes have been identified as the products of biosynthetic
evolution of NO, suggested as a means of NO storage, and found to be responsible for
several important biological functions.ii Among those, dinitrosyl iron complexes (DNICs)
have also been identified in the macrophages responsible for killing invading
microorganisms by releasing NO.iii They are also found to be cytotoxic to E-coli
bacteriaiv and some of them exhibit anti-tumor activities.v These discoveries have
stimulated intense interest in isolating and studying metal-nitrosyl complexes, especially
those that mimic the structures of biologically active non-heme iron nitrosyl complexes,vi
and those capable of releasing NO for pharmacological and medicinal applications.vii
Bidentate phosphine ligands are commonly used in organic chemistry as
activators for catalyst complexes. Coordination complexes of triphenylphosphine,
chelating agents such as 1,2-bis(diphenylphosphino)ethane or 1,1-bis(diphenyl-
phosphino)methane are generally used as homogeneous catalysts for a wide range of
reactions. These catalysts frequently make use of transition metals like palladium and
iridium.viii A few examples of phosphine disubstituted nitrosyl complexes have been
previously synthesized as early as 1962 with some recent examples reported as well.ix
However the vast majority of these compounds are not chelated species and
crystallographic data was not provided, nor was a detailed analysis of the compounds
done.x
4

In continuing our interest in the preparation of novel dinitrosyl iron complexes,
we have reported chelating compounds using N,N-chelating ligands,xi and dinuclear
dinitrosyliron complexes linked by bis(diphenyl)phosphine ligands.xii In this paper,
reactions using Fe(NO)2(CO)2 and a series of bis(diphenylphosphine) chelating ligands
ranging in sizes and the presence or absence of π-bonds in the carbon backbones, such as
1,2-bis(diphenyl-phosphino)benzene, (DPPBz), 1,3-bis(diphenylphosphino)propane,
(DPPP), and cis-1,2-bis(diphenylphosphino)ethylene, (DPPEt) were investigated. From
these reactions, three chelated products with the general formula [(DPPX)Fe(NO)2],
where DPPX = C6H4[P(C6H5)2]2, (C6H5)2PCH2CH2CH2P(C6H5)2, and
(C6H5)2PCH=CHP(C6H5)2 were isolated and characterized using methods such as FT-IR,
UV-Vis, and NMR spectroscopies as well as X-Ray crystallography and cyclic
voltammetry. Quantum mechanical calculations using density functional theory were
performed on these compounds and their model compounds using two different methods.
The calculated results were found to be consistent with the experimental observations.
Materials and Methods:
All solvents were purchased from Aldrich in anhydrous form or obtained dry and
air free from a MBRAUN MB-SPS dispenser (Manual Solvent Purification System) and
stored in small amounts over 3 or 4 Å molecular sieves. Pentacarbonyliron(0), 1,2-
bis(diphenylphosphino)benzene (DPPBz), cis-1,2-bis(diphenylphosphino)ethylene
(DPPEt), and 1,3-bis(diphenylphosphino)propane (DPPP) were all purchased from Sigma
Aldrich and used without further purification. Dicarbonyldinitrosyliron(0) was
synthesized using a modified method of Seel’s procedure.xiii,xiv Because the
5

Fe(CO)2(NO)2 is very volatile and tends to decompose over time, it was stored in
individual 1 mL aliquots as a 1:1 solution with dichloromethane or THF for a later use.
Synthesis of Fe(NO)2(DPPBz) (1)
1,2-Bis(diphenylphosphino)benzene (DPPBz) ligand (0.405 g, 0.907 mmol) was
dissolved in a minimal amount of THF (no more than 7 mL) in a glovebox at room
temperature with stirring in a 25 mL round bottom flask fitted with a rubber septum.
Then 0.2 mL (0.0907 mmol) of Fe(CO)2(NO)2/THF solution (1:1 ratio) was introduced
into the flask via syringe. The solution was allowed to stir at room temperature overnight
to produce the monosubstituted product, Fe(NO)2(CO)(DPPBz). After stirring, the
reaction mixture was taken out of the glovebox and gently heated at 50°C for 24 hours
with the reaction progress monitored through FT-IR spectroscopy. The solution was
transferred back into the glovebox after completion of the double substitution. Large dark
purple crystals of X-ray quality ranging in size from 0.5 mm3 to 3 mm3 were grown by
allowing the solvent to slowly evaporate over several days through a vent needle. The
crystals and solid powder from the bottom of the flask were filtered and washed with
three 0.5 mL aliquots of cool toluene. Yield: 0.343g, 72%; IR (in THF) ν(NO) (cm-1): 1719
1
and 1675. NMR (δ-ppm): H, 7.170 (s) and 7.323 (d, J = 7.8 Hz) (C6H5), 7.546 (m)
(C6H4); 13C, 126.0 (s, p-C6H5), 128.8 (s) and 129.6 (s) (o- or m-C6H5), 133.4 (t, JC-P = 6.4
Hz) and 129.1 (t, JC-P = 5.1 Hz) (ipso-C6H5 or C6H4), 138.4 (s) and 130.6 (s) (o- or m-
C6H4). Elemental analysis: Expt (calc): C, 51.47% (64.08%); H, 4.10% (4.30%); N,
3.72% (4.98%).
6

Synthesis of Fe(NO)2(DPPP) (2)
Synthesis of the chelated species was done by adding 0.374 g (0.907 mmol) of
1,3-bis(diphenylphosphino)propane (DPPP) to 10 mL THF at room temperature with
stirring under a nitrogen atmosphere in a 50 mL round bottom flask with a rubber septum.
Then 0.2 mL (0.0907 mmol) of a 1:1 Fe(CO)2(NO)2/THF solution was introduced into
the flask via syringe. The solution was allowed to stir at room temperature overnight to
obtain the monosubstituted product, Fe(NO)2(CO)(DPPP). After stirring, the reaction
mixture was transferred out of the glovebox and gently heated at 50°C. IR analysis of the
liquid phase showed a sharp decrease in the intensity of the remaining CO peak after 24
hours and very little change in the intensity after 36 hours of heating. The mixture was
transferred into the glovebox where the solvent was allowed to slowly evaporate. Small
needle-shaped red crystals of X-ray quality ranging in size from 0.5mm3 to 1mm3 were
grown by allowing the solvent to slowly evaporate over several days. The crystals and
solid powder from the bottom of the flask were filtered and washed with three 0.5 mL
aliquots of cool toluene. Yield: 0.302g, 63%; IR (in CH2Cl2) ν(NO) (cm-1): 1708 and 1660.
1
NMR (δ-ppm): H, 2.027 (m) and 2.427 (m, JH-P = 10.8 Hz, JH-H = 7.8 Hz)
(CH2CH2CH2), 7.272 (m) and 7.455 (m) (20 H, C6H5); 13C, 20.3 (s) (-CH2-), 28.7 (d, JC-P
= 2.5 Hz) (CH2-R-CH2), 135.4 (doublet of doublets, JC-P(1) = 33.2 Hz, JC-P(2) = 2.7 Hz)
(ipso-C6H5), 128.5 (d, JC-P = 12.7 Hz) and 132.2 (d, JC-P = 10.0 Hz) (o- or m-C6H5), 129.8
(s) (p-C6H5). Elemental Analysis: Expt (calc) C, 61.70% (61.38%); H, 5.01% (4.96%); N,
5.19% (5.30%).
Synthesis of Fe(NO)2(DPPEt) (3)
7

In order to synthesize the chelated product, 0.359 g (0.907 mmol) of (cis-1,2-
bis(diphenylphosphino)ethylene) DPPEt was dissolved in 10 mL THF at room
temperature under a nitrogen atmosphere in a 50 mL round bottom flask with a rubber
septum. Then, 0.2 mL (0.0907 mmol) of a 1:1 Fe(CO)2(NO)2/THF solution was
introduced into the flask via syringe. The solution was allowed to sit at room temperature
overnight. IR analysis showed the presence of only one CO stretching peak, indicating
the loss of one of the carbonyls and substitution with one of the phosphine groups. The
flask was transferred out of the glovebox and gently heated at 50°C for 24 hrs. IR
analysis of the liquid phase showed a sharp decrease in the intensity of the remaining CO
peak after 24 hrs and very little change in the intensity after 36 hrs of heating. The
mixture was transferred inside the glovebox where the solvent was allowed to slowly
evaporate. Large dark purple plate crystals of X-ray quality ranging in size from 1 mm3 to
5 mm3 were grown by allowing the solvent to slowly evaporate over several days. The
crystals and solid powder from the bottom of the flask were filtered and washed with
three 0.5 mL aliquots of cool toluene. Yield: 0.353g, 76%; IR (in CH2Cl2) ν(NO) (cm-1):
1723 and 1674. NMR (δ-ppm): 1H, 7.436 (m) (20H, C6H5), 7.738 (2H, CH=CH); 13C,
134.5 (d, JC-P = 21.2 Hz) (ipso-C6H5), 132.2 (t, JC-P = 6.4 Hz) and 128.9 (t, JC-P = 5.1 Hz)
(o- or m-C6H5), 130.3 (s) (p-C6H5), 148.8 (t, JC-P = 33.6 Hz) (CH=CH). Elemental
analysis: Expt (calc): C, 59.74% (60.96%); H, 4.40% (4.33%); N, 5.09% (5.47%).
Spectroscopic Measurements:
All Solid and liquid Fourier Transform Infrared (FT-IR) spectra were recorded on
a Nicolet FT-IR iS10 spectrometer using the Smart MIracle ZnSe ATR Accessory with a
8

ZnSe crystal or a Smart OMNI-Transmission accessory with KBr pellets in the range of
400-4000 cm-1. Air sensitive solids and liquids were prepared in the glovebox and
scanned using the Smart OMNI-Transmission accessory under a dry nitrogen
environment. 1H and 13C NMR spectra were recorded on a Bruker 300 MHz NMR
spectrometer, using tetramethylsilane as an internal standard. NMR samples were
prepared in the glove box under nitrogen by dissolving an appropriate amount of product
in 0.5 ml of deuterated solvent and filtering any undissolved sample through glass wool.
The UV-Visible absorption spectra were recorded on a Shimadzu UV-2100U
spectrophotometer. The samples were prepared inside a glovebox under a nitrogen
atmosphere using Quartz airtight cells with anhydrous dichloromethane as a solvent.
X-ray crystallography:
X-ray crystallographic data for complexes 1, 2 and 3 were collected from a single
crystal sample mounted on a glass fiber. The instrument used for the diffraction data
collection was a Bruker SMART APEX II diffractometer equipped with a fine focus,
2.0kW sealed tube X-ray source (MoKα radiation, λ = 0.71073Å) operating at 50kV and
30mA. Data collection was carried out by APEX II software. SAINT-PLUS was used to
integrate the intensity of reflections, while SADABS was used for scaling and absorption
correction. The structures were solved using the heavy atom method. All non-hydrogen
atoms were refined anisotropically. Computations were performed using SHELXTL and
final full-matrix refinements were against F2. xv
9

Electrochemistry:
The electrochemical studies were performed on a CH Instruments electrochemical
analyzer 660A. A three-electrode system consisting of a platinum working electrode, a
platinum wire counter electrode and an Ag/Ag+ reference electrode was used. The
reference electrode was separated from the bulk of the solution by a fritted-glass bridge
filled with a solvent/supporting electrolyte mixture. Cyclic voltammograms were
recorded in a CH2Cl2 solution of 0.1 M tetrabutylammonium perchlorate (TBAP) with a
sample concentration of approximately 0.001 M and a scan rate from 100 mVs-1 to 900
mVs-1. A background spectrum was taken for each aliquot of electrolyte solution to
ensure that the solvent and electrolyte had not become contaminated. The cyclic
voltammogram of ferrocene was recorded each day before a set of experiments in order
to counteract any drift from the reference electrode. The potential values are reported vs
+
ferrocene/ferrocenium ion; the E1/2[Fe(Cp)2/Fe(Cp)2 ] under our experimental conditions is
0.132 V.
Theoretical calculations:
The equilibrium geometries of the three compounds were calculated using density
functional theory (DFT) with Becke’s three-parameter exchange potential and the Lee-
Yang-Parr correlation functional (B3LYP) along with the 6-31+G(d) basis set.xvi In order
to confirm the convergence of the geometries with the basis set, additional B3LYP
calculations using the aug-cc-pVDZ basis set as well as the 6-31+G(d) basis set were
performed on the model systems representing the three compounds. In these model
10

systems, methyl groups were used to replace the phenyl groups in the original
compounds.
Harmonic frequency calculations were carried out for the three compounds and
the model systems following geometry optimization in our calculations. All harmonic
frequencies obtained were found to be real, confirming each of the equilibrium
geometries calculated corresponded to a minimum on the potential energy surface.
Additionally, some of the calculated harmonic frequencies, particularly the NO stretch
frequencies, can be directly compared with the experimental frequencies. All calculations
were carried out using the GAUSSIAN 09 computational chemistry program.xvii
Results and Discussion
Reactions:
Compounds 1-3 were synthesized by first generating the corresponding
monosubstituted compounds at room temperature and then heating the reaction mixture at
50°C for 24 – 36 hrs. Attempts to synthesize the chelated species through immediate
heating of the reaction mixture without first producing the monosubstituted species lead
to decomposition and overall loss of product. The general synthesis of the
monosubstituted and chelated compounds is depicted in Scheme 1.
It has been reported that the reaction kinetics for the substitution of the first
carbonyl can follow pseudo-first order kinetics,xviii therefore the rate of the reaction will
be dependent on the concentration of the nucleophile. By introducing Fe(CO)2(NO)2 to
the phosphine ligand in high concentrations (using a minimal amount of solvent to
dissolve the ligand before injecting the starting material) the reaction yields nearly pure
11

monosubstituted product in as little as one hour. The formation of the monosubstituted
compounds was monitored using FT-IR spectroscopy in order to confirm that a shorter
reaction time was possible under the new synthesis conditions. It was observed that the
first CO group was almost entirely replaced within approximately 40 minutes.
The second carbonyl substitution occurs over a longer period of time when
compared with the first. The chelation reaction itself depends on the orientation of the
second phosphine group being correct to replace the second carbonyl. IR analysis of the
solutions of the monosubstituted products shows that a small amount of the chelated
species can form without the application of heat. This behavior was only observed in
complexes 1 and 3 where the conjugated carbon bridge locks the phosphine in the correct
orientation. This behavior is demonstrated by the presence of several downshifted NO
peaks in the IR spectra of the compounds. The chelation process for these two complexes
is only dependent on the rotation of the iron-phosphorus and phosphorus-carbon bonds of
the ligand, and not on the rotation of the long saturated carbon chain present in the ligand
of compound 2.
The compounds 1-3 are insoluble in nonpolar organic solvents such as toluene
and pentane but are moderately soluble in halogenated solvents like dichloromethane and
chloroform. The monosubstituted complexes are highly air sensitive and will decompose
rapidly, forming a black sticky solid upon being exposed to air. The chelated compounds
appear to be air stable for up to 12 hours in the crystalline form. In a solution of
dichloromethane, compounds 1 and 3 are stable for several days while compound 2
decomposes over 8 hours.
12

1H and 13C NMR spectroscopy showed well-resolved spectra and all complexes
were found to be relatively pure, with the exception of compound 1, where the NMR
showed some solvent residues present from the synthesis. Coupling of carbons and
protons to the phosphorous atoms at up to three bonds away was observed, and all
carbons and protons have been identified in the spectra. Elemental analysis for 2 and 3
matches well with the theoretical one, except for compound 1, possibly due to the
presence of solvent as shown in NMR.
Infrared Spectroscopy
The synthesis of both monosubstituted and chelated compounds was qualitatively
monitored through the use of IR spectroscopy to ensure the completion of the reactions.
The loss of carbonyl stretching peaks is easily noted due to their usual intensity and
predictable location. The starting material Fe(CO)2(NO)2 exhibits two nitrosyl stretches
at 1797 and 1725 cm-1 and two carbonyl stretches at 2080 and 2018 cm-1. Upon
replacement of one of the carbonyls with a phosphine ligand, the νNO stretches shift to
lower values in the region of 1750 to 1715 cm-1, resulting from an increase in back-
donation of electron density from the d-orbital of the metal center into the π* orbital of
the NO group. The replacement of the second carbonyl by the phosphine group causes a
shift of both NO stretching frequencies to even lower wavenumbers in the range of 1650
to 1750 cm-1, as listed in Table 1.
Interestingly, the disubstituted compounds 1 and 3 both shift to wavenumbers in
the range of 1725-1720 cm-1 and 1675-1670 cm-1, while compound 2 and a similar
compound using 1,2-bis(diphenylphosphino)ethane (DPPE)xix exhibit an even greater
13

shift in the range of 1708-1707 cm-1 and 1660-1657 cm-1. This indicates that compound
2 and the DPPE complex both have N-O bonds that are weakened more than those of
compounds 1 and 3 after substitution. This trend illustrates the fact that compound 2 and
the DPPE compound have phosphorus atoms that act as stronger nucleophiles during the
substitution reaction and are overall stronger σ-donors. Conversely, the ligands in 1 and
3 act as weaker nucleophiles, likely due to the presence of C=C π-bonds between the two
phosphorus atoms. It is possible that the double bond allows for partial redistribution of
the electron density from the two phosphorus atoms, leading to a lower amount of
electron density available for donation from the phosphorus to the iron center and
consequently, less back-donation from the metal to the NO group. This is also confirmed
by the X-ray crystallographic analysis and the theoretical calculations, which will be
discussed later.
UV-Vis Spectroscopy
The UV-Visible absorption spectra for complexes 1-3 show strong absorption
peaks in the ultraviolet region around 248 nm, a shoulder from 275 to 320 nm, and a
weaker broad peak from 475 to 600 nm. The UV region of the spectrum is several orders
of magnitude more intense than the visible region. The peak at 248 nm has been assigned
to π – π* transitions from the four phenyl rings of the ligand because the same absorption
was also observed in the ligands alone. The shoulder from 275 to 320 nm in the UV
region are possibly due to a metal to ligand charge transfer (MLCT), which is strong
because the absorption is both spin-allowed and Laporte-allowed. The broad peaks from
475 to 600 nm in the visible region (Figure 1) are much weaker and could be attributed to
14

d-d transitions, characteristic of transition metal compounds. Molar absorptivity
calculations were done by taking the absorbance at the highest point of the broad peak at
505 nm at several concentrations. Plots of absorption vs concentration yielded the molar
absorptivity coefficients of 184.6 M−1 cm−1, 166.4 M−1 cm−1 and 241.4 M−1 cm−1 for
complexes 1, 2 and 3, respectively. These values are typical for d-d transitions of
transition metal complexes with tetrahedral geometry.
X-ray Crystallography
Single crystals of X-ray quality ranging in size from 1 mm3 to 5 mm3 were grown
by allowing the solvent to slowly evaporate and the X-ray crystal structures of the
chelated compounds 1-3 were determined by mounting a suitable single crystal on a glass
fiber. Crystallographic collection and refinement parameters can be found in Table 2.
Selected bond lengths and bond angles can be found in Table 3 and Table 4, respectively.
All three compounds crystallized in the monoclinic crystal system, with space groups of
P2(1)/c for compounds 1 and 2 and C2/c for 3. The iron centers on all three crystalline
compounds show the distorted tetrahedral geometry typical of DNIC species. The X-ray
crystal structures are shown in Figures 2-4.
Compounds 1 and 3 consist of distorted five-membered rings formed by the Fe
atom, two phosphorus atoms and two carbon atoms. This is to be expected since the sp2
hybridization of the double bonded carbon backbones holds the phosphorus atoms at a
rigid angle of 120°. Because of this geometry, both compounds exhibit P-Fe-P bonds
that are less than 90° and Fe-P bond distances that are longer than a normal C-C single
bond. Compound 2 consists of a six-membered ring containing three carbon atoms, two
15

phosphorus atoms and the iron atom. By comparison, the two phosphorus atoms of the
ligand in compound 2 are not held in a rigid orientation as the carbon chain connecting
the phosphorus atoms contains three single-bonded carbons instead of two carbons with
double bonds as in 1 and 3. This allows for compound 2 to have a larger P-Fe-P bond
angle of 94.45(3)°. However, the P-Fe bond lengths are similar to those of compounds 1
and 3 at ~2.2 Å. Compound 2 also exhibits a six-membered ring in a distorted chair
conformation due to the contracted P-Fe-P bond angle.
Dinitrosyl iron complexes are frequently described as having “attracto”
conformations where the N-Fe-N bond angles are generally less than or equal to 130° but
are greater than the O-Fe-O bond angles.xx This conformation leads to the two oxygen
atoms being slightly bent towards each other. The attracto conformation is usually
observed in first-row transition metal complexes containing a good π-acceptor ligand.
Compound 2 exhibits a N-Fe-N bond angle of 126.42(11)°, which is greater than the O-
Fe-O bond angle of 124.47˚. This compound exhibits the expected attracto conformation
where the oxygen atoms are bent towards each other. This fact is reinforced upon
considering that the Fe-N-O bond angles of the compound are 175-176°, which is close to
linear, but still differs by 4-5°. On the contrary, compounds 1 and 3 have N-Fe-N bond
angles of ~129° {129.46(8)° and 129.80(12)°}, O-Fe-O bond angles of ~130-131° and
nearly linear Fe-N-O bond angles of 178-179°. This indicates that these two complexes
are of the rarer “repulso” conformation. An exaggerated molecular illustration of the
attracto and repulso conformations for these compounds is shown in Figure 5.
Typically, the repulso conformation is more common for the second-row and
third-row transition metals.xxi For example, the N-M-N and O-M-O bond angles reported
16

for M(NO)2(PPh3)2 complex are 139.2° and 142.7° when M = Ruxxii and 139.1° and
140.6° when M = Os.xxiii This is also true for the cationic species, [M(NO)2(PPh3)2]+,
where the reported N-M-N and O-M-O bond angles are 157.5° and 173.7° for Rh and
154.2° and 167.5° for Ir, respectively.xxiv
In order to determine whether steric or electronic effects dictate the different
conformations between attracto and repulso, another complex similar to compound 2 was
examined. The compound Fe(NO)2(DPPE),19 which differs from compound 2 by only
one carbon atom, has a P2(1)/c space group similar to compound 1. The N-O bond
distances for this structure (1.185 and 1.200 Å) are very similar to those of the chelated
structures synthesized in this work, while the Fe-N bonds (1.658 and 1.681 Å) are similar
as well. This indicates that the DPPE ligand experiences back-donation in a similar
manner to 1, 2 and 3. One difference lies in the angles of the iron atom and the NO
groups. The N-Fe-N bond angle for the DPPE complex is 125.4° which is similar to
compound 2. However the O-Fe-O bond angle for the DPPE complex is given as 124.0°.
These angles indicate that this complex is of the attracto conformation similar to 2,
despite the fact that it also contains a 5-membered ring like compounds 1 and 3. The
repulso conformations exhibited by compounds 1 and 3 are quite rare. Only one similar
compound has been crystallized in the repulso conformation; Co(NO)2(DPPE) with a N-
Co-N bond angle of 131.7° and an O-Co-O bond angle of 135.9°.xxv
When examining the IR stretching frequencies for compounds 2 and
Fe(NO)2(DPPE), it can be seen that they both have similar values. Thus both phosphine
ligands donate electron density in a manner that is more intense than the ligands
containing π-bonds. Because compounds 1, 3 and the DPPE complex all have five-
17

membered rings, but 1 and 3 are repulso while the DPPE is attracto, the conformational
preference cannot be due to a steric effect. The difference between the attracto and
repulso forms must then be an electronic effect due to 1 and 3 containing C=C π bonds,
while 2 and the DPPE complex do not.
The bond distances of the NO groups reveal yet more information about the
nature of the phosphorus-iron bonds. The Fe-N bond distances at ~1.65 Å are contracted
while the N-O bond distances of ~1.2 Å are slightly lengthened. This indicates that there
is a significant amount of back-donation of electron density from the iron atom into the
π*-orbitals of the NO ligands. This supports the idea that the NO bonds are being
weakened as shown by the IR spectral data mentioned earlier.
Cyclic Voltammetry
The redox behaviors of complexes 1-3 were studied through cyclic voltammetry
as shown in Figure 6. The CV of each of the compounds was first compared with that of
their respective ligand to identify whether any redox reactions are coming specifically
from the ligand, and not the iron center. Compound 1 exhibits two quasi-reversible
oxidations at 0.101 and 0.186 V. The separation between the two couples is small (0.086
V), but can still be distinguished visually. The presence of the two peaks can be verified
by viewing the second derivative spectrum. Compound 2 exhibits a different behavior
with two distinct and separate oxidation couples at 0.019 and 0.342 V. The separation
between the two oxidation couples is quite large (0.323 V). The CV of compound 3
appears to have only one peak but the peak is broad and, by examining the second
derivative CV, it can be confirmed that there are actually two separate oxidation couples
18

at 0.121 and 0.184 V. Table 5 summarizes the E1/2 values for the three compounds.
The separations between the two oxidation couples E°1/2 (1) and E°1/2 (2) for
complexes 1, 2, and 3 are 0.085, 0.323 and 0.063 V, respectively. It is found that the
separation between redox couples for complexes 1 and 3 are much smaller than that of
compound 2, which is likely due to the structure of the chelating ligand, as compounds 1
and 3 both have C=C π bonds present while complex 2 does not. It is proposed that when
the first oxidation occurs on the metal center, the positive charge can be delocalized
through the C=C double bond into the chelating ligand itself, which makes the iron
charge less than +1. As a result, the second oxidation would be easier, which renders the
second oxidation E°1/2 (2) to occur at a potential closer to the first oxidation E°1/2 (1). As a
consequence, the difference between the first oxidation and second oxidation is smaller.
On the contrary, compound 2 lacks the π system and, therefore, cannot delocalize the
charge away from the iron center. This means the iron has a full +1 charge, and removing
a second electron from the iron center is more difficult. This is reflected by the second
oxidation, which occurs at higher potential, resulting the larger separation between
E°1/2(2) and E°1/2(1) values.
The CVs of variable scan rates from 100 mV to 900 mV were recorded for each
compound and a plot of intensity of the current vs. the square root of the scan rate (Figure
7) shows a linear relationship indicating that all processes are diffusion-controlled single-
electron redox couples.
Theoretical Calculations:
19

B3LYP density functional theory calculations using the 6-31+G* basis set were
carried out for each of the three compounds. In addition, B3LYP calculations using 6-
31+G* and aug-cc-pVDZ basis sets were carried out for the model compounds, in which
each of the phenyl groups on the phosphorus is replaced by a methyl group. The
calculated bond lengths and bond angles matched well with the experimental data as
shown in Tables 3 and 4. For compounds 1, 2 and 3, the calculated average for the bond
lengths of N-O using the 6-31+G* method are 1.1885, 1.1887 and 1.1862 Å, which are
very close to the experimental averages, at 1.1875(2), 1.193(3) and 1.1865(3) Å,
respectively. Using model compounds, the corresponding bond lengths are 1.1879,
1.1900 and 1.1876 Å for the 6-31+G* method and 1.1840, 1.1860 and 1.1836 Å for the
aug-cc-pVDZ method, respectively. For Fe-N bond lengths, the average calculated values
using the 6-31+G* method are 1.6376, 1.6373 and 1.6377 Å, which are systematically
shorter than the average for the experimental ones {1.6446(17), 1.6495(2) and 1.651(2)
Å} respectively. Using model compounds, the corresponding bond lengths are 1.6399,
1.6394 and 1.6396 Å for the 6-31+G* method and 1.6403, 1.6396 and 1.6400 Å for the
aug-cc-pVDZ method, respectively, which are closer to the experimental values. The
experimental average bond length for Fe-P is 2.2183(5), 2.2419(7) and 2.2297(7) Å, for
compounds 1, 2 and 3, while the corresponding calculated averages are 2.2380, 2.2430,
and 2.2391 Å for the 6-31+G* method, 2.2113, 2.2149, and 2.2172 Å for the 6-
31+G*(model), and 2.2324, 2.2404 and 2.2395 Å for the aug-cc-pVDZ(model) method,
respectively. The aug-cc-pVDZ(model) produced best fit for experimental data, where
the average bond length of Fe-P is longer for 2 than that of 1 and 3.
It is observed that the IR stretching frequencies for the NO group are 10-15 cm-1
20

lower for complex 2 than for complexes 1 and 3, indicating stronger back donation from
the DPPP ligand than DPPBz and DPPEt ligands. Our calculated harmonic frequencies,
as shown in Table 1, are in good agreement with the trend in the IR frequencies. It should
be noted that all calculated harmonic frequencies are about 5% higher than the
corresponding experimental frequencies, which is typical because the unharmonicity
correction would be expected to fill such a difference. The trend in the IR frequencies for
the NO group is also consistent with the X-ray crystal structure data, where the average
bond lengths for N-O is 1.193(3) Å for 2, which is longer than that of 1 and 3 {1.1875(2)
and 1.1865(3) Å, respectively}, because the back-donation weakens the N-O bond. The
calculated data using all three methods consistently showed that the bond lengths of the
N-O group are longer for compound 2 than that of compound 1 and 3.
The calculated bond angles that dictates the local geometry of the metal center,
such as P-Fe-P and P(1,2)-Fe-N(1,2) match well with the experimental data. The
experimental average of the Fe-N-O angles are 178.56° and 179.05° for compound 1 and
3, which are more linear than that of compound 2 (176.18°). The calculated average Fe-
N-O angles using 6-31+G* are 177.43°, 176.83° and 176.80° for 1, 2 and 3, respectively.
When using the model compounds, the average bond angles are 179.21°, 179.43° and
179.04° for the 6-31+G* method and 178.70°, 178.81° and 178.51° for the aug-cc-pVDZ
method for 1, 2 and 3, respectively. The calculated bond angles of N-Fe-N and O-Fe-O
are quite similar for the three compounds and only attracto conformations are predicted,
21

while experimentally 2 exhibits attracto but 1 and 3 possess repulso conformations.
The highest occupied molecular orbital (HOMO) - lowest unoccupied molecular
orbital (LUMO) energy separations are obtained from the calculations as well. The
HOMO-LUMO gaps for compounds 1, 2 and 3 are 3.736, 4.060, and 3.669 eV,
respectively. It is interesting to note that these HOMO-LUMO gaps are very large, which
makes the compounds very stable with respect to reduction, thus may account for the fact
that only oxidations are observed in the electrochemical study.
It is also noted that the compounds 1 and 3 are similar, while compound 2 is
significantly larger. It is consistent with the fact that the observed charge transfer bands at
~320 nm in the UV-Vis spectra {Figure 1(a)} have similar intensity for compounds 1 and
3, but are much stronger for compound 2. From the experimental and calculated bond
lengths and FT-IR data, it is clear that the DPPP ligand acts as better donor to the metal,
thus the metal has more electrons available for back-donation for compound 2 than
compounds 1 and 3. It has been reported that depleting electron density from the metal
center through metal to ligand back-bonding results in widening the HOMO-LUMO
gap,xxvi which is consistent with our results. Moreover, it is known that the HOMO-
LUMO gap is lowered by the presence of conjugated π-systems in organic molecules.xxvii
Therefore, the double bond in between the two phosphorus atoms may account for the
lower HOMO-LUMO gap for compounds 1 and 3.
CONCLUSIONS
We have demonstrated that dinitrosyl iron complexes containing chelated
structures of phosphine ligands including 1,2-bis(diphenylphosphino)benzene, cis-1,2-
22

bis(diphenylphosphino)ethylene and 1,3-bis(diphenylphosphino)propane, can be
synthesized and characterized by IR, UV-Vis, NMR, X-ray and CV methods.
Infrared spectroscopy was used to qualitatively monitor the reaction progress. The
reaction times were found to be dependent on the presence of coordinating and non-
coordinating solvents. The first substitution occurred over one hour in THF vs. 24 hours
in dichloromethane, indicating that THF may act as a hard base and stabilizes the loss of
the first carbonyl group.
Single crystals of compounds 1, 2 and 3 were obtained and the X-ray
crystallographic data shows they crystallize in the monoclinic crystal system with
P2(1)/c, P2(1)/c and C2/c space groups and the unit cells contain 4, 4 and 8 molecules,
respectively. The average Fe-N-O bond angles for the three compounds are all close to
linear with compounds 1 and 3 at 178.5° while compound 2 at 175.5°, which is smaller.
Compounds 1 and 3 possess five-membered rings and have P-Fe-P bond angles less than
90° while compound 2 has an angle of 94° and contains a six-membered ring. Complex 2
exhibits the attracto state while 1 and 3 exhibit the repulso conformations, which is rare
for the first row transition metals. The presence of π-bonds on the carbon bridge of the
bis(diphenyl)phosphine ligands plays a significant role in the observed conformation,
while steric hindrance is eliminated as a cause.
The UV-Visible absorption spectra for complexes 1-3 show strong absorptions at
248 nm, a shoulder from 275 to 320 nm, and a weaker broad peak from 475 to 600 nm,
which are assigned to π – π* transitions from the four phenyl rings of the ligand, metal to
ligand charge transfer (MLCT), and d-d transitions, respectively. Molar absorptivity
coefficients for the visible spectra are calculated to be 184.6 M−1 cm−1, 166.4 M−1 cm−1
23

and 241.4 M−1 cm−1 for complexes 1, 2 and 3, respectively, which are typical for d-d
transitions of transition metal complexes with tetrahedral geometry.
The redox behaviors of all three compounds were studied through cyclic
voltammetry. All three compounds in the series exhibit two quasi-reversible oxidation
peaks. The E°1/2 values for compounds 1, 2 and 3 are 0.101 and 0.186 V, 0.121 and 0.184,
and 0.019 and 0.342V, respectively. The separation between the two oxidation couples is
much smaller for compounds 1 and 3 when compared with that of compound 2, which
has been attributed to the presence of π-bonds in the bisphosphine chelating ligands,
allowing for charge delocalization upon the first oxidation occurring.
DFT calculations using the B3LYP functional and the 6-31+G* basis set were
performed on all three compounds. In addition, B3LYP calculations using the 6-31+G*
and aug-cc-pVDZ basis sets were carried out for the model compounds, in which each of
the phenyl groups on the phosphorus is replaced by a methyl group. The calculated bond
lengths and angles match well with the experimental data and the calculated IR stretching
frequencies are also consistent with experimental observations. The calculated HOMO-
LUMO gaps for compounds 1, 2 and 3 are 3.736, 4.060, and 3.669 eV, respectively. The
higher HOMO-LUMO gap for compound 2 is consistent with more electron back-
donation from the metal to the NO ligands, while the lower HOMO-LUMO gaps for
compounds 1 and 3 are attributed to the conjugated π-systems on the DPPBz and DPPEt
ligands. These large HOMO-LUMO gaps are also consistent with the lack of the usual
reductions in electrochemical studies.
24

Appendix A. Supplementary data
CCDC 965700, 965701 and 965702 contain the supplementary crystallographic
data for compounds 1, 2 and 3. These data sets can be obtained free of charge via
http://www.ccdc.cam.ac.uk/conts/retrieving.html, or from the Cambridge
Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: (+44)
1223-336-033; or e-mail: deposit@ccdc.cam.ac.uk.
Acknowledgements:
We wish to thank the National Institute of Health (NIH) MBRS SCORE Program
(Grant #1SC3GM092301-01) for financial support.
25

Table 1. List of Experimental and Calculated IR Stretching Frequencies for CO and NO.
CO (cm-1)
ν
NO (cm-1)
References
Compounds
ν
Fe(CO)2(NO)2
2080 and 2018
1797 and 1725
1760, 1719a
1719, 1675a
1814, 1763d
1759, 1716a
1708, 1660b
1807, 1756d
1758, 1716a
1723, 1674b
1825, 1760 d
1764, 1722
1724, 1678
1707, 1657
This work
This work
Fe(NO)2(CO)(DPPBe)
Fe(NO)2(DPPBe) (1)
1994
This work
This work
This work
This work
This work
This work
This work
This work
10
-
-
Fe(NO)2(CO)(DPPPr)
2002
Fe(NO)2(DPPPr) (2)
-
-
Fe(NO)2(CO)(DPPEt)
2001
Fe(NO)2(DPPEt) (3)
-
-
Fe(NO)2(CO)P(C6H5)3
Fe(NO)2[P(C6H5)3]2
Fe(NO)2(DPPE)
2009
5
-
10
-
(a) - in THF at room temperature (b) - in dichloromethane at room temperature (c) Solid KBr
Pellet. (d) Calculated frequencies using the B3LYP/6-31+G(d) method.
26

Table 2. Crystallographic Collection and Refinement Parameters for 1, 2 and 3.
Parameter
Compound 1
Compound 2
Compound 3
Empirical Formula C30 H24 Fe N2 O2 P2 C27 H26 Fe N2 O2 P2 C26 H22 Fe N2 O2 P2
Molecular Weight
Description
Size mm3
Temperature , K
Wavelength, Å
Crystal System
Space group
a, Å
562.30
Dark Purple Prism
528.29
Red needle
512.25
Dark Purple Plate
3
3
3
4.0 x 0.4 x 0.3 mm 3.0 x 3.0 x 0.4 mm 5.0 x 3.0 x 2.0 mm
230(2) K
0.71073 Å
Monoclinic
P2(1)/c
9.9462(11) Å
15.4195(17) Å
17.828(2) Å
90°
296(2) K
0.71073 Å
Monoclinic
P2(1)/c
9.4281(3) Å
15.9709(6) Å
16.7774(7) Å
90°
296(2) K
0.71073 Å
Monoclinic
C2/c
17.9383(6) Å
14.6440(6) Å
18.8183(6) Å
90°
b, Å
c, Å
α, deg
β, deg
92.218(5)°
90°
93.437(2)°
90°
95.914(2)°
90°
γ, deg
Volume Å3
3
3
3
2732.1(5) Å
4
2521.72(16) Å
4
4917.0(3) Å
8
Z
Calc. Density
3
3
3
1.367 Mg/m
1.392 Mg/m
1.384 Mg/m
mg/m3
Abs. Coeff. mm-1
F(000)
-1
-1
-1
0.699 mm
1160
0.752 mm
1096
0.769 mm
2112
θ Range deg.
1.75 to 25.15°
-10<=h<=11
-18<=k<=18
-21<=l<=20
31834
1.76 to 25.05°
-11<=h<=11
-18<=k<=13
-18<=l<=19
11208
1.80 to 25.01°
-21<=h<=17
-14<=k<=17
-22<=l<=22
8472
Index Ranges
No. Refl. Coll.
No. Indep. Refl.
4883 [R(int) =
0.0211]
4442 [R(int) =
0.0209]
4315 [R(int) =
0.0253]
Data / restr. / param 4883 / 0 / 430
4442 / 0 / 411
4315 / 0 / 386
Goodness-of-fit on
1.000
1.063
1.006
2
F
Final R indices
[I>2sigma(I)]
R1 = 0.0291, wR2
= 0.0843
R1 = 0.0348, wR2
= 0.0842
R1 = 0.0367, wR2
= 0.0944
R1 = 0.0357, wR2
= 0.0909
R1 = 0.0433, wR2
= 0.0905
R1 = 0.0513, wR2
= 0.1035
R indices (all data)
0.431 and -0.181
0.338 and -0.233
0.469 and -0.219
Largest diff. peak
and hole
-3
-3
-3
e.Å
e.Å
e.Å
28

Table 3: List of Selected Experimental and Calculated Bond Lengths (Å) for Compounds
1, 2 and 3.
Bond
Fe(1)-N(1) Expt
6-31+G*
6-31+G*(model)
Method
Compound 1 Compound 2 Compound 3
1.6485(16)
1.6373
1.652(2)
1.6379
1.6397
1.6398
1.190(3)
1.1881
1.1903
1.1864
1.647(2)
1.6366
1.6390
1.6394
1.196(3)
1.1892
1.1896
1.1856
2.2501(7)
2.2364
2.2194
2.2404
2.2336(7)
2.2496
2.2194
2.2404
1.650(2)
1.6389
1.6396
1.6400
1.186(3)
1.1817
1.1875
1.1836
1.652(2)
1.6365
1.6396
1.6400
1.187(3)
1.1907
1.1875
1.1836
2.2228(7)
2.2407
2.2172
2.2395
2.2297(7)
2.2375
2.2172
2.2395
1.6399
aug-cc-pVDZ(model) 1.6403
N(1)-O(1)
Expt
1.182(2)
1.1910
1.1879
6-31+G*
6-31+G*(model)
aug-cc-pVDZ(model) 1.1840
Fe(1)-N(2) Expt
6-31+G*
6-31+G*(model)
1.6404(17)
1.6368
1.6399
aug-cc-pVDZ(model) 1.6403
N(2)-O(2)
Expt
1.193(2)
1.1859
1.1879
6-31+G*
6-31+G*(model)
aug-cc-pVDZ(model) 1.1840
Fe(1)-P(1) Expt
6-31+G*
6-31+G*(model)
2.2180(5)
2.2378
2.2112
aug-cc-pVDZ(model) 2.2324
Fe(1)-P(2) Expt 2.2185(5)
6-31+G* 2.2382
6-31+G*(model) 2.2114
aug-cc-pVDZ(model) 2.2324
29

Table 4: List of Selected Experimental and Calculated Bond Angles (deg) for
Compounds 1, 2 and 3.
Bonds
Compound 1 Compound 2
Compound 3
129.80(12)
122.47
126.31
125.68
86.28(3)
87.44
N(1)-Fe(1)-N(2) Expt
129.46(8)
123.76
126.42(11)
123.89
127.59
126.81
94.45(3)
96.33
6-31+G*
6-31+G*(model)
126.44
aug-cc-pVDZ(model) 125.82
P(1)-Fe(1)-P(2)
Expt
86.67(2)
86.17
6-31+G*
6-31+G*(model)
88.08
96.62
87.56
aug-cc-pVDZ(model) 87.54
95.70
87.02
O(1)-Fe(1)-O(2) Expt
6-31+G*
6-31+G*(model)
131.17
121.61
125.78
124.47
121.46
127.11
125.81
176.68
176.36
179.74
179.27
175.67
176.97
179.12
178.35
106.49(8)
103.29
106.11
106.18
100.48(7)
130.31
119.80
125.50
124.44
179.50
177.52
179.04
178.51
178.6
aug-cc-pVDZ(model) 124.73
Fe(1)-N(1)-O(1) Expt
178.90
176.92
179.22
6-31+G*
6-31+G*(model)
aug-cc-pVDZ(model) 178.73
Fe(1)-N(2)-O(2) Expt
178.23
177.94
179.20
6-31+G*
176.08
179.04
178.51
112.33(8)
107.39
109.03
109.34
107.64(8)
6-31+G*(model)
aug-cc-pVDZ(model) 178.66
N(1)-Fe(1)-P(1) Expt 109.33(5)
6-31+G* 110.62
6-31+G*(model) 108.86
aug-cc-pVDZ(model) 109.12
N(1)-Fe(1)-P(2) Expt 107.61(5)
30

6-31+G*
110.63
108.84
109.31
106.11
106.17
112.48(8)
110.52
108.04
108.79
111.32(8)
109.82
108.04
108.77
107.37
109.03
109.33
102.08(8)
113.67
109.03
109.33
110.07(9)
112.84
109.03
109.33
6-31+G*(model)
aug-cc-pVDZ(model) 109.12
N(2)-Fe(1)-P(1) Expt 106.66(6)
6-31+G* 109.64
6-31+G*(model) 108.94
aug-cc-pVDZ(model) 109.27
N(2)-Fe(1)-P(2) Expt 108.75(6)
6-31+G* 109.64
6-31+G*(model) 108.95
aug-cc-pVDZ(model) 109.27
Table 5: List of Redox Potentials for Compounds 1, 2 and 3, their Cathodes and Anodes
Peak Separations.
Compounds
∆E(1)
∆E(2)
E°1/2(1)
(V) a
0.132
E°1/2(2)
(V) a
-
[E°1/2(2) - E°1/2(1)]
(mV)b
(mV)b
(mV)
Ferrocene
89
73
63
69
-
0.101
0.019
0.121
0.186
0.342
0.184
64
79
65
85
323
63
1
2
3
a Values corrected using the E1/2 value of ferrocene/ferrocenium as a standard.
b ∆E(1) = Epa – Epc for E°1/2(1); ∆E(2) = Epa –Epc for E°1/2(2).
31