Fine tuning of the oxidation locus, and electron transfer, in nickel complexes of pro-radical ligands

Article abstract of DOI:10.1002/chem.200500915

A large number of complexes of the first-row transition metals with non-innocent ligands has been characterized in the last few years. The localization of the oxidation site in such complexes can lead to discrepancies when electrons can be removed either

Full text of DOI:10.1002/chem.200500915

FULL PAPER
DOI: 10.1002/chem.200500915
Fine Tuning of the Oxidation Locus, and Electron Transfer, in Nickel
Complexes of Pro-Radical Ligands
Olaf Rotthaus,[a] Olivier Jarjayes,*[a] Fabrice Thomas,*[a] Christian Philouze,[a]
Carlos Perez Del Valle,[a] Eric Saint-Aman,[b] and Jean-Louis Pierre[a]
Abstract: A large number of com-
plexes of the first-row transition metals
with non-innocent ligands has been
characterized in the last few years. The
localization of the oxidation site in
such complexes can lead to discrepan-
cies when electrons can be removed
either from the metal center (leading
to an M(n+1)+ closed-shell ligand) or
from the ligand (leading to an Mn+
open-shell ligand). The influence of the
ligand field on the oxidation site in
square-planar nickel complexes of
redox-active ligands is explored herein.
The tetradentate ligands employed
herein incorporate two di-tert-butylphe-
nolate (pro-phenoxyl) moieties and
one orthophenylenediamine spacer.
The links between the spacer and both
phenolates are either two imines
([Ni(L1)]), two amidates ([Ni(L3)]2ꢀ),
or one amidate and one imine
([Ni(L2)]ꢀ). The structure of each nick-
el(ii) complex is presented. In the non-
coordinating solvent CH2Cl2, the one-
electron-oxidized forms are ligand–rad-
ical species with a contribution from a
singly occupied d orbital of the nickel.
In the presence of an exogenous
ligand, such as pyridine, a NiIII closed-
shell ligand form is favored: axial liga-
tion, which stabilizes the trivalent
nickel in its octahedral geometry, indu-
ces an electron transfer from the met-
al(ii) center to the radical ligand. The
affinity of pyridine for the phenoxyl-
nickel(ii) species is correlated to the N-
donor ability of the linkers.
Keywords: electron
transfer
·
nickel · radicals · Schiff bases ·
tautomerism
Introduction
Consequently, a large number of complexes of the first-row
transition metals with noninnocent ligands have been char-
acterized.[2] These ligands, in their diamagnetic form, mainly
incorporate phenoxyl or (imino)semiquinonate coordinating
groups. With the aim of modeling the active site of galactose
oxidase, we have recently prepared several phenoxyl cop-
per(ii) complexes.[3] In particular, we have shown that the
phenolatocopper(ii) complexes of H2L1 and other salen-
type ligands can undergo a one-electron oxidation into mag-
netically interacting phenoxyl copper(ii) complexes.[4,5] The
localization of the oxidation site in such complexes could,
however, lead to discrepancies as electrons can be removed
either from the metal center (leading to an M(n+1)+ closed-
shell ligand) or from the ligand (leading to an Mn+ open-
shell ligand), and this situation could become even more
complicated when valence tautomerism exists between these
two forms, as in nickel porphyrins:[7] in this case, the two
forms—the M(n+1)+ closed-shell porphyrin and the Mn+
open-shell porphyrin—are in equilibrium.
A fascinating class of metalloproteins use free radicals as co-
factors to promote oxidation reactions.[1] Galactose oxidase,
which belongs to this class, nicely illustrates the high degree
of sophistication that can be reached in biocatalysis when a
radical and a metal work in synergy. Several biomimetic ap-
proaches have been developed by chemists during the last
decade to get insights into the metal–radical interaction.
[a] Dr. O. Rotthaus, Dr. O. Jarjayes, Dr. F. Thomas, Dr. C. Philouze,
Dr. C. Perez Del Valle, Prof. Dr. J.-L. Pierre
Laboratoire de Chimie BiomimØtique, LEDSS, UMR CNRS 5616
UniversitØ J. Fourier
BP 53, 38041, Grenoble Cedex 9 (France)
Fax : (+33)4-7651-4836
E-mail: Olivier.Jarjayes@ujf-grenoble.fr
Fabrice.Thomas@ujf-grenoble.fr
[b] Prof. Dr. E. Saint-Aman
LEOPR, UMR CNRS 5630
UniversitØ J. Fourier, BP 53
38041 Grenoble Cedex 9 (France)
An understanding of the properties of such systems is of
major interest not only from a theoretical point of view but
also for the design of molecular electronic devices and
switches.[8] This has, for example, recently stimulated the
Supporting information for this article is available on the WWW
under http://www.chemeurj.org/ or from the author.
Chem. Eur. J. 2006, 12, 2293 – 2302
ꢀ 2006 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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Table 1. Crystallographic data for [Ni(L1)], [Ni(L2)]ꢀ, and [Ni(L3)]2ꢀ
[Ni(L1)] [Ni(L2)]ꢀ [Ni(L3)]2ꢀ
C36H46N2NiO2 C42H61N3NiO3 C68H116N4NiO4
.
formula
M
crystal system
space group
a [Š]
b [Š]
c [Š]
a [8]
b [8]
597.47
714.66
monoclinic
P21/a
12.108(2)
26.330(5)
12.901(3)
90
95.59
90
4093(1)
4
1112.39
triclinic
P1
triclinic
¯
¯
P1
9.190(2)
12.981(3)
14.386(4)
107.98(2)
93.46(2)
95.56(2)
1617.3(8)
2
13.753(4)
15.770(6)
16.969(7)
107.46(2)
110.33(3)
80.06(3)
3282(3)
2
g [8]
V [Š3]
Z
T/K
150
1.227
6.33
293
1.160
5.13
150
1.125
3.43
1calcd [gcmꢀ3
]
m [cmꢀ1
]
monochromator
radiation
graphite
MoKa
graphite
MoKa
graphite
MoKa
wavelength [Š]
reflections collected
independent reflections
0.71073
41422
9384
(0.14400)
7811
[I>2s(I)]
0.0557
0.0786
0.71073
14325
6515
(0.06301)
4603
[I>2s(I)]
0.0538
0.0715
0.71073
61054
16559
(0.16874)
9191
[I>2s(I)]
0.0765
0.1125
synthesis of nickel complexes with dioxolene ligands that ex-
hibit such valence tautomerism.[9]
(Rint
)
In slightly distorted square-planar nickel complexes, the
energy levels of the potentially redox-active orbitals—p
from the ligand or d from the metal—are very similar: by
shifting one above the other (modulation of the ligand
field), a metal-based or ligand-based redox process can be
favored. With this goal, we present in this paper the one-
electron oxidative chemistry of a series of nickel(ii) com-
plexes of pro-phenoxyl ligands possessing two di-tert-butyl-
phenolates and two N-chelating groups: two imines for
H2L1, one imine and one amidate for H3L2, and two ami-
dates for H4L3. The ligand field is shown to dramatically
affect the oxidation locus in these compounds.
observed reflections
R
Rw
Results and Discussion
Figure 1. X-ray crystal structure of [Ni(L1)] (1) shown with 30% thermal
ellipsoids.
Preparation of the ligands: H2L1 was obtained by condensa-
tion of o-phenylenediamine with 3,5-di-tert-butylbenzalde-
hyde. Coupling 3,5-di-tert-butylsalicyl chloride and 2-amino-
1-(3,5-di-tert-butyl-2-hydroxybenzimino)benzene in CH2Cl2
affords H3L2. H4L3 was obtained by coupling o-phenylenedi-
amine with 3,5-di-tert-butylsalicyl chloride.
Structures and spectroscopic properties of the nickel(ii)
complexes: The three ligands provide an N2O2 coordination
sphere for a single nickel ion. Mixing one equivalent of Ni-
(OAc)2·4H2O, two equivalents of NEt3, and H2L1 in ethanol
affords [Ni(L1)] (1). Similarly, the reaction of H3L2 with two
equivalents of Ni(OAc)2·4H2O and four equivalents of NEt3
(containing one equivalent of nBu4N+OHꢀ) in ethanol gave
the low-spin complex [Ni(L3)]ꢀ (2ꢀ). [Ni(L3)]2ꢀ (32ꢀ) was ob-
tained by adding 1.2 equivalents of Ni(OAc)2·4H2O to H4L3
in DMF in the presence of four equivalents of nBu4N+OHꢀ.
The X-ray crystal structures (Table 1) of 1, 2ꢀ, and 32ꢀ are
depicted in Figures 1–3, respectively, and selected bond
lengths and angles are reported in Table 2. The metal ion is
coordinated to two imine or amidate nitrogens (N1, N2) and
two phenolate oxygens (O1, O2). The N2O2Ni coordination
Figure 2. X-ray crystal structure of [Ni(L2)]ꢀ (2) shown with 30% thermal
ellipsoids.
polyhedron is essentially square planar but is slightly distort-
ed towards tetrahedral.
ꢀ
ꢀ
ꢀ
ꢀ
The Ni O1, Ni O2, Ni N1, and Ni N2 bond lengths in 1
(Figure 1) are roughly similar (1.854(1), 1.850(1), 1.852(1)
and 1.857(1) Š, respectively). This behavior contrasts with
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Chem. Eur. J. 2006, 12, 2293 – 2302

Nickel Complexes of Pro-Radical Ligands
FULL PAPER
ꢀ
ꢀ
are similar to those in 1. The Ni O2 and Ni N2 distances in
the o-hydroxybenzamidate moiety are 1.843(2) and 1.871(2),
respectively. The mean deviation from the least-squares
plane defined by the atoms N1, N2, O1, O2, and Ni is
0.058 Š, with a dihedral angle between the two O-Ni-N
planes of 1778. The phenyl ring C8–C13 is more distorted
than in 1, as reflected by a mean deviation from the least-
squares plane of 0.0147 Š. The N1 and N2 atoms are dis-
placed 0.0987 and ꢀ0.0025 Š, respectively, from this plane,
with a torsion angle around the N1-C8-C13-N2 bond of 0.98.
Figure 3. X-ray crystal structure of [Ni(L3)]2ꢀ (3) shown with 30% ther-
mal ellipsoids.
2ꢀ
ꢀ
ꢀ
The Ni O1 and Ni O2 bonds in 3 (1.853(3) and
ꢀ
1.845(3) Š, respectively; Figure 3) are shorter than the Ni
ꢀ
N1 and Ni N2 bonds (1.876(3) and 1.867(3) Š, respective-
ly). The mean deviation from the least-squares plane de-
fined by the atoms N1, N2, O1, O2, and Ni is 0.065 Š, a
value that is intermediate between those of 1 and 2. The di-
Table 2. Selected bond lengths [Š] and angles [8] for [Ni(L1)], [Ni(L2)]ꢀ,
and [Ni(L3)]2ꢀ
.
[Ni(L1)]
1.850(1)
ꢀ
ꢀ
ꢀ
ꢀ
ꢀ
hedral angle between the two O Ni N planes is 1798, as for
1. The phenyl ring C8–C13 deviates significantly from pla-
narity (mean deviation of 0.0318 Š), with the nitrogen
atoms N1 and N2 being displaced 0.2371 and ꢀ0.2218 Š, re-
spectively, from this plane.
Ni O1
ꢀ
Ni N2
1.854(1)
1.857(1)
1.420(2)
1.308(2)
85.9(1)
Ni O2
ꢀ
Ni N1
1.852(1)
1.303(2)
1.419(3)
94.7(1)
ꢀ
O1 C3
C1 N1
O1-Ni-N1
N1-Ni-N2
1.303(2)
1.305(2)
94.6(1)
85.5(1)
O2 C16
ꢀ
ꢀ
ꢀ
C1 C2
C14 C15
ꢀ
N2 C14
O2-Ni-N2
O1-Ni-N2
O1-Ni-O2
O2-Ni-N1
174.7(1)
173.2(1)
[Ni(L2)]
The UV/Vis spectrum of 1 in CH2Cl2 (Figure 4, Table 3) is
dominated by intense absorptions at 383 (e=
29600 mꢀ1 cmꢀ1), 470 (sh; e=7000), 492 (e=9200), and a low
intensity shoulder at around 600 nm (e<200). The 490-nm
band, which is absent in the optical spectrum of the free
ꢀ
ꢀ
ꢀ
Ni O1
ꢀ
Ni N2
1.856(2)
1.871(2)
1.422(5)
1.338(4)
85.1(1)
Ni O2
1.843(2)
1.301(4)
1.289(4)
93.8(1)
86.4(1)
Ni N1
1.857(2)
1.307(4)
1.479(4)
95.1(1)
ꢀ
ꢀ
O1 C3
O2 C16
ꢀ
ꢀ
ꢀ
C1 C2
C1 N1
C14 C15
ꢀ
N2 C14
O1-Ni-N1
N1-Ni-N2
O2-Ni-N2
O1-Ni-N2
O1-Ni-O2
O2-Ni-N1
173.8(1)
177.0(1)
[Ni(L3)]
ꢀ
ꢀ
ꢀ
Ni O1
ꢀ
Ni N2
1.853(3)
1.867(3)
1.491(5)
1.354(5)
83.3(1)
Ni O2
1.845(3)
1.310(5)
1.357(5)
95.4(1)
86.7(1)
Ni N1
1.876(3)
1.308(5)
1.493(5)
95.1(1)
ꢀ
ꢀ
O1 C3
O2 C16
ꢀ
ꢀ
ꢀ
C1 C2
C1 N1
C14 C15
ꢀ
N2 C14
O1-Ni-N1
N1-Ni-N2
O2-Ni-N2
O1-Ni-N2
O1-Ni-O2
O2-Ni-N1
175.4(1)
174.1(1)
that of [Cu(L1)] (the nickel atom has been replaced by a
copper ion),[10] which exhibits Cu O1 and Cu O2 bonds
(1.895(2) and 1.894(2) Š, respectively) that are shorter than
ꢀ
ꢀ
Figure 4. Electronic spectra of 0.05mm solutions of [Ni(L1)] (solid lines)
and electrogenerated [Ni(L1)]+ (dashed lines) in CH2Cl2 (+0.005m
TBAP). The dotted lines represent [Ni(L1)]+ in the presence of 5mm pyr-
idine. T=233 K, l=1.000 cm.
ꢀ
ꢀ
the Cu N1 and Cu N2 ones (1.941(2) and 1.945(2) Š, re-
spectively). The mean deviation from the least-squares
plane defined by the atoms N1, N2, O1, O2, and Ni is
0.0786 Š, with a dihedral angle between the O1-Ni-N1 and
O2-Ni-N2 planes of 1798. The
nitrogen atoms N1 and N2 are
Table 3. Electronic properties of the nickel complexes in CH2Cl2 solution.
displaced 0.0903 and 0.0255 Š,
respectively, from the plane de-
fined by the atoms C8, C9, C10,
C11, C12, and C13 (mean devi-
ation of 0.0100 Š) and the tor-
sion angle N1-C8-C13-N2 is
0.58.
Complexes
lmax [nm] (e [mꢀ1 cmꢀ1])[a]
[Ni(L1)]
383 (29600), 492 (9200)
[Ni(L1)]+
370 (21700), 413 (12200), 458 sh (7560), 780 br (2420), 936 br (2480)
325 (32200), 342 sh (28600), 362 sh (21500), 468 (16100)
375 (9620), 417 (9900), 445 (10000), 544 (3060), 918 (2900)
325 (10600), 366 (10400), 427 (9400)
345 sh (11600), 404 (9600), 432 (8600), 463 sh (6680), 922 br (2600)
308 (20000), 433 (8600)
329 (17600), 374 (21400), 393 sh (17400), 501 (1000)
342 (20000), 372 sh (15600), 414 sh (7600), 795 (2400)
326 (30400), 457 (6400), 655 br (2800)
[Ni(L1)(Py)2]+
[Ni(L1)]2+
[Ni(L2)]ꢀ
[Ni(L2)]
[Ni(L2)(Py)2]
[Ni(L3)]2ꢀ
[Ni(L3)]ꢀ
In the nickel complex 2ꢀ of
the unsymmetrical tetradentate
2
ꢀ
[Ni(L3)(Py)2]ꢀ
[Ni(L3)]
ligand H3L (Figure 2), the Ni
ꢀ
O1 and Ni N1 distances of the
337 (6100), 383 sh (4600), 476 (1600), 618 br (1600)
salicylidene moiety (1.856(2)
[a] UV/Vis spectra recorded at 233 K in the presence of 0.005m TBAP (CH2Cl2 +0.1m TBAP solutions electro-
and 1.857(2) Š, respectively) chemically oxidized and diluted twenty times); sh=shoulder, br=broad.
Chem. Eur. J. 2006, 12, 2293 – 2302
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2295

F. Thomas, O. Jarjayes et al.
ligand, corresponds to a charge-transfer (CT) transition. As
expected, this absorption is red-shifted (by 50 nm) and is
less intense (by a factor of 2.5) than the CT transition of the
previously described [Cu(L1)] complex.[5,10,11] The shorter-
wavelength features are mainly associated with intraligand
p–p* transitions, while the longer-wavelength bands corre-
spond to d–d transitions.
The electronic spectrum of 2ꢀ (Figure 5, Table 3) exhibits
absorptions at 325 (e=10600 mꢀ1 cmꢀ1), 366 (e=10400), 427
(e=9400), 500 (shoulder; e=4000), and 580 nm (shoulder;
Figure 6. Electronic spectra of 0.05mm solutions of [Ni(L3)]2ꢀ (solid lines)
and electrogenerated [Ni(L3)]ꢀ (dashed lines) in CH2Cl2 (+0.005m
TBAP). The dotted lines represent [Ni(L3)]ꢀ in the presence of 25mm
pyridine. T=233 K, l=1.000 cm.
trolyte) by cyclic voltammetry (CV), differential pulse vol-
tammetry (DPV), and rotating disc electrode (RDE) vol-
tammetry. All potentials are referenced to the ferrocenium/
ferrocene couple.
In CH2Cl2, the CV curves of 1 and 2ꢀ (Figure 7, Table 5)
2
exhibit two reversible redox waves (at E1/21 =0.582, E1/2
=
Figure 5. Electronic spectra of 0.05mm solutions of [Ni(L2)]ꢀ (solid lines)
and electrogenerated [Ni(L2)] (dashed lines) in CH2Cl2 (+0.005m
TBAP). The dotted lines represent [Ni(L2)] in the presence of 5mm pyri-
dine. T=233 K, l=1.000 cm.
0.802 V and E1/21 =0.150, E1/22 =0.806 V versus Fc+/Fc for
e=500) whose assignment is as follows: the higher-wave-
length feature corresponds to d–d transitions, the CT transi-
tions are observed in the 370–500 nm range, while the short-
er-wavelength absorptions are attributed to intraligand p–p*
transitions.
The visible spectrum of [Ni(L3)]2ꢀ (Figure 6, Table 3) in
CH2Cl2 reveals transitions at 329 (shoulder; e=
17600 mꢀ1 cmꢀ1), 374 (e=21400), 501 (e=1000), and 550 nm
(shoulder; e=400). The two latter bands are attributed to
d–d transitions, while the former correspond to p–p* and
CT transitions. Similar transitions are found in the UV/Vis
spectra of 1, 2ꢀ, and 32ꢀ in CH3OH (Table 4).
Figure 7. CV curves of 1mm CH2Cl2 solutions (+0.1m TBAP) of: a)
[Ni(L1)], b) [Ni(L2)]ꢀ, and c) [Ni(L3)]2ꢀ. Scan rate: 0.1 Vsꢀ1; T=298 K.
The potentials are referenced against Fc+/Fc.
Electrochemistry of the nickel(ii) complexes: The electro-
chemical behavior of the nickel complexes was studied in
CH2Cl2 and CH3OH (with 0.1m TBAP as supporting elec-
[Ni(L1)] (1) and [Ni(L2)]ꢀ (2ꢀ), respectively), while three
waves are observed for [Ni(L3)]2ꢀ (32ꢀ; at E1/21 =ꢀ0.254,
E1/22 =0.306, and E1/23 =0.706 V versus Fc+/Fc). All of these
correspond to one-electron pro-
cesses, as judged by coulometry
Table 4. Spectroscopic properties of nickel complexes in MeOH solution.
and RDE voltammetry. The E1/
Complexes
lmax [nm] (e [mꢀ1 cmꢀ1])[a]
1
2
oxidation potentials are re-
[Ni(L1)]
381 (29800), 486 (9600)
markably dependent on the co-
[Ni(L1)]+
317 (29800), 334 sh (25500), 355 sh (21300), 460 (14400)
317 (27700), 334 sh (24500), 356 sh (19700), 457 (13300)
314 (17600), 446 (6800), 930 br (800)
321 (10800), 370 (9600), 408 (8200), 481 (3800)
313 (20800), 432 (9000)
ordinating nitrogens, and thus
on the global charge of the
complex: the higher value is ob-
served for 1 (two imines), the
lower for 32ꢀ (two amidates),
while an intermediate potential
is obtained for 2ꢀ. In the pres-
[Ni(L1)(Py)2]+
[Ni(L1)]2+
[Ni(L2)]ꢀ
[Ni(L2)]
[Ni(L2)(Py)2]
[Ni(L3)]2ꢀ
[Ni(L3)]ꢀ
306 (21000), 411 (8600)
334 (13600), 363 (11400), 388 sh (8200), 515 (600)
327 (15400), 361 sh (10800), 821 (2600)
316 (18800), 466 (2800), 645 br (1000)
[Ni(L3)(Py)2]ꢀ
1
ence of pyridine, the E1/2 and
[a] UV/Vis spectra recorded at 233 K in the presence of 0.005m TBAP (CH2Cl2 +0.1m TBAP solutions electro-
chemically oxidized and diluted twenty times); sh=shoulder, br=broad.
2
E1/2 values are shifted towards
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Nickel Complexes of Pro-Radical Ligands
Table 5. Redox properties of the nickel complexes
FULL PAPER
Complex
[Ni(L1)]
Solvent
E1/2 [V][a]
CH2Cl2
CH2Cl2 +5 equivof Py
MeOH
CH2Cl2
CH2Cl2 +50 equivof Py
MeOH
CH2Cl2
CH2Cl2 +1000 equivof Py
MeOH
0.582
0.574
0.466
0.150
0.102
0.802
0.758
0.670
0.806
0.61
0.514
0.306
0.074
0.290
[Ni(L2)]ꢀ
[Ni(L3)]2ꢀ
0.234
ꢀ0.254
ꢀ0.322
ꢀ0.054
0.706
0.626
[a] Potential values given relative to the Fc/Fc+ reference electrode, T=
298 K, complex concentrations are 1mm. E1/2 values were obtained from
DPV measurements by adding half of the pulse amplitude to the Ep
value. The confidence level is ꢂ0.005 V.
Figure 8. X-band EPR spectra of CH2Cl2 solutions (+0.1m TBAP) of: a)
1mm [Ni(L1)]+ at 100 K, b) 10mm [Ni(L2)] at 100 K, c) 1mm [Ni(L3)]ꢀ at
100 K, and d) 1mm [Ni(L1)]+ at 260 K. Microwave frequency: 9.42 GHz;
power: 20 mW; modulation frequency: 100 kHz; amplitude: 0.4 mT.
the cathodic region, which suggests that the oxidized forms
are stabilized in the presence of an exogenous N-donor
ligand (Table 5).
cule. The temperature-dependence of the EPR spectra of 1+
shows that the change of shape (NiIII to radical) occurs at
around 170–190 K, which is the melting point of the sol-
vent.[15] A similar behavior has recently been reported by
Yamauchi et al. and interpreted as a valence tautomerism
governed by the temperature.[16] An alternative explanation
could be relaxation effects: as the metal relaxes faster than
the radical, only the radical contribution is detected in
liquid CH2Cl2. Because none of these arguments is particu-
larly strong, we still entertain both possibilities.
The CV curves of 1, 2ꢀ, and 32ꢀ in the coordinating sol-
vent CH3OH show waves that are less reversible than those
obtained in CH2Cl2; the oxidation potentials are slightly dis-
1
placed. A cathodic shift of E1/2 is observed for the Schiff-
base complex 1, while it is anodic for 2ꢀ and 32ꢀ: there is
1
therefore no simple correlation between E1/2 and the struc-
ture of the complex in CH3OH.
One-electron-oxidized complexes in CH2Cl2: The one-elec-
tron-oxidized complexes 1+, 2, and 3ꢀ were generated elec-
trochemically in CH2Cl2 (with 0.1m TBAP as supporting
electrolyte) at 233 K under argon.
Both 2 and 3ꢀ are EPR-silent at 233 K,[17] although lower-
ing the temperature to 100 K results in the appearance of a
rhombic (S=1/2) signal (Figure 8, Table 6), with a moderate
g-tensor anisotropy (g=2.09, 2.03, 1.95 and 2.048, 2.020,
1.999 for [Ni(L2)] and [Ni(L3)]ꢀ, respectively) and gav values
of 2.022 and 2.015, respectively. These are lower than any
reported for NiIII complexes[14] or high-spin nickel(ii) mag-
netically coupled to a radical.[18] On the other hand, both
the gav values and the g-tensor anisotropy are higher than
those of phenoxyl radicals, thus suggesting that 2 and 3ꢀ are
ligand radicals with contribution from the singly occupied d
orbital of NiIII.[19] The EPR spectrum of 2 does not match
either that of 1+ (i.e., a salicylidene(phenoxyl)nickel com-
plex) or that of 3ꢀ (i.e., an o-hydroxybenzamidate-
The electronic spectra of 1+, 2, and 3ꢀ exhibit remarkable
bands in the 800–1000 nm region (Figures 4–6). Based on
their high intensity (1000<e<3400 mꢀ1 cmꢀ1), they corre-
spond to CT transitions. Similar bands, although less intense,
were observed in the phenoxyl radical complex [Cu(L1)]+.[5]
Additional features are observed in the 400–450 nm region,
where the p–p* transitions of phenoxyl radicals are expect-
ed to be found.[12] These data thus suggest a ligand-centered,
rather than metal-centered, oxidation process at 233 K. The
one-electron-oxidized species are highly stable in solution at
298 K, with less than 5% degradation within 10 h for 1+ and
2 and less than 3% degradation within six days for 3ꢀ.
The X-band EPR spectrum of 1+ at 260 K in CH2Cl2
(+0.1m TBAP) exhibits an isotropic and unresolved (S=1/2)
signal centered at giso =2.034 (peak-to-peak line width of
3.1 mT, Figure 8). This giso value is intermediate between
that obtained for the phenoxyl zinc(ii) analogue [Zn(L1)]+[13]
and that of typical NiIII complexes (roughly 2.13–2.17).[14] 1+
is thus best formulated as a ligand radical, with partial de-
localization of the unpaired electron onto the orbitals of the
nickel atom. The EPR spectrum of 1+ in CH2Cl2 (+0.1m
TBAP) at 100 K contrasts sharply with that recorded at
260 K. It consists of a rhombic (S=1/2) signal, with g values
at 2.270, 2.220, and 2.022 (gav =2.17) that are typical of NiIII
complexes in a low-spin configuration. The pattern of g
values (gx ꢁgy >gz ꢁge) is indicative of a dz2 ground state
where the z-axis is perpendicular to the plane of the mole-
(phenoxyl)nickel
complex),
which suggests a contribution
from the two resonance struc-
tures shown in Scheme 1 to the
electronic distribution in 1+, 2,
and 3ꢀ. These results were con-
Scheme 1. Resonance
struc-
firmed by DFT calculations on tures for [Ni(L3)]ꢀ.
3ꢀ performed at the B3LYP
level (Figure 9): the calculated
SOMO is mainly developed on the ligand, with contribution
of a nickel dxy, dxz, or dyz orbital. As expected, the ligand
SOMO is equally developed on both phenoxyl moieties,
thus indicating the delocalization of the unpaired electron as
depicted in Scheme 1.
The three complexes 1+, 2, and 3ꢀ thus exhibit a ligand
radical character, with contribution of a nickel orbital.
Chem. Eur. J. 2006, 12, 2293 – 2302
ꢀ 2006 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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2297

F. Thomas, O. Jarjayes et al.
Table 6. EPR parameters of the one-electron-oxidized nickel complexes.
Complex
Solvent[a]
g values
A values
gxx
gyy
gzz
gav
Axx
Ayy
Azz
[Ni(L1)]+
CH2Cl2
2.270
2.200
2.287
2.205
2.048
2.218
2.309
2.218
2.086
2.227
2.328
2.223
2.220
2.168
2.228
2.170
2.020
2.181
2.246
2.183
2.017
2.203
2.270
2.195
2.022
2.027
2.023
2.026
1.999
2.026
2.020
2.027
1.941
2.024
2.018
2.024
2.171
2.132
2.179
2.134
2.022
2.142
2.192
2.143
2.015
2.151
2.205
2.147
CH2Cl2 +Py[b]
MeOH
1.65
1.65
1.60
1.55
1.55
1.55
1.65
1.65
1.55
1.55
1.55
1.55
2.14
2.14
2.10
2.10
2.04
2.04
MeOH+Py[b]
CH2Cl2
[Ni(L2)]
CH2Cl2 +Py[b]
MeOH
MeOH+Py[b]
CH2Cl2
[Ni(L3)]ꢀ
CH2Cl2 +Py[b]
MeOH
MeOH+Py[b]
[a] At 100 K in the presence of 0.1m TBAP. [b] 0.1m pyridine.
Figure 10. X-band EPR spectra of 1mm CH2Cl2 solutions (+0.1m TBAP
+ 5000 equivalents of pyridine) of: a) [Ni(L1)]+ at 233 K, b) [Ni(L1)]+ at
100 K, c) [Ni(L2)] at 100 K, and d) [Ni(L3)]ꢀ at 100 K. Solid lines repre-
sent experimental spectra and dotted lines are simulations using the pa-
rameters given in Table 6. Microwave frequency: 9.42 GHz; power:
20 mW; modulated frequency: 100 kHz; amplitude 0.1 mT.
Figure 9. Optimized structure and calculated SOMO for [Ni(L3)]ꢀ. Calcu-
lations performed at the B3LYP level using the Lanl2dz basis set.
tions at wavelengths higher than 600 nm and medium-inten-
sity CT transitions in the range 400–600 nm (Table 3).
Among the surveyed compounds, 3ꢀ has the lowest affini-
ty for pyridine, with a logb value at 298 K for the equilibri-
um shown in Equation (1) of only 2.41ꢂ0.06 mꢀ2; logb is
6.71ꢂ0.09 mꢀ2 for 1+.[20] The stronger donor capacity of the
amido N-donors thus lowers the affinity of the complexes
for axial ligands.[21]
One-electron-oxidized complexes in CH2Cl2/pyridine and
CH3OH: Increasing amounts of pyridine were added to
CH2Cl2 solutions of the radical complexes 1+, 2, and 3ꢀ at
298 K and dramatic color changes were observed: the
orange solution of 1+ turns brown, as do the green solutions
of 2 and 3ꢀ. The evolution of the electronic spectra during
titration follows a similar trend, namely the disappearance
of the 900–1000-nm absorption.
For 1+, a broad absorption, assigned to NiIII d–d transi-
tions, appears around 700–900 nm, whereas the shorter-
wavelength absorption at around 400–600 nm is attributed
to CT transitions. This NiIII, rather than ligand radical, state
was confirmed by EPR spectroscopy at 298 K, which reveals
an isotropic (S=1/2) signal centered at a giso value of 2.15
(Figure 10a).[14]
No other absorptions could be observed above 600 nm in
the electronic spectrum of 2 in the presence of pyridine, al-
though medium-intensity bands, attributed to CT, are pres-
ent at 560 (shoulder; e=2340 mꢀ1 cmꢀ1) and 468 nm (e=
18460). Similarly, the visible spectrum of 3ꢀ in the presence
of 10000 equivalents of pyridine exhibits NiIII d–d transi-
2
½NiðLފ þ 2 Py Ð ½NiðLÞðPyÞ2Š b ¼ ½NiðLÞðPyÞ2Š=ð½NiðLފ ꢃ ½PyŠ Þ
ð1Þ
Lowering the temperature from 298 to 233 K results in an
increase of logb by three units for 3ꢀ and two units for 1+.
As the color of the pyridine adduct is different from that of
the radical complex, a thermochromism is observed for 1+
and 3ꢀ in the presence of one and 100 equivalents of pyri-
dine, respectively. This thermochromism is more marked for
3ꢀ than for 1+, in agreement with the difference of logb
(see Supporting Information).
2298
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Chem. Eur. J. 2006, 12, 2293 – 2302

Nickel Complexes of Pro-Radical Ligands
FULL PAPER
The EPR spectra of the pyridine adducts of 1+, 2, and 3ꢀ
in CH2Cl2 at 100 K are shown in Figure 10. A low-spin (S=
1/2) NiIII signal, with hyperfine splitting in each of the three
g components, is systematically observed. The pattern of g
values is consistent with an octahedral geometry around the
metal. In addition, the presence of one well-resolved quintu-
plet in the high-field component, as well as two less-resolved
quintuplets in the low-field component, implies the exis-
tence of the six-coordinate NiIII adducts [Ni(L1)(Py)2]+ (4+),
[Ni(L2)(Py)2] (5), and [Ni(L3)(Py)2]ꢀ (6ꢀ) with two equiva-
lent pyridines axially bonded (2”2IN +1 transitions; IN =
1).[22] Replacement of the imine nitrogens by amidate indu-
ces both a lowering of Azz and an increase of gxx and gyy due
to a weakening of the axial bonds.[23] This explanation is
consistent with the measured pyridine binding, which fol-
lows the order 1+ >2>3ꢀ.
octahedral NiIII moiety at 100 K (with two CH3OH mole-
cules axially bonded), and a radical UV/Vis spectrum at
233 K, similar to 1+ in CH2Cl2. This may be explained by
the weaker donor ability of CH3OH (compared to pyridine)
and the lower affinity of axial ligands for this complex,
which does not favor an octahedral geometry.
Conclusion
The three ligands H2L1, H3L2, and H4L3 afford the slightly
distorted square-planar nickel(ii) complexes 1, 2ꢀ, and 32ꢀ,
respectively, if no exogenous ligands are present. The corre-
sponding one-electron-oxidized[24] species in CH2Cl2 are
highly stable delocalized radicals with a contribution from
the nickel orbital. Strong amido N-donors, which are known
to stabilize high-valent metals, do not promote a simple
metal-centered oxidation process. The energy levels of the
potentially redox-active orbitals—p from the ligand, or d
from the metal—thus remain almost the same in 2 and 3ꢀ,
as seen by the respective EPR spectra at 100 K (complex
1+, which contains imino N-donors, exhibits a more marked
nickel(iii) character). The shift of one orbital above the
other, that is, an intramolecular electron transfer (valence
tautomerism), can be promoted by changing the geometry
(from square planar to octahedral) around the metal center.
The corresponding one-electron-oxidized complexes are
nickel(iii) phenolate complexes.
The electronic spectra of the oxidized species 1+ are
roughly similar in CH3OH and CH2Cl2/pyridine. The same
behavior is observed for 2, which suggests that the oxidation
products in CH3OH contain NiIII. Addition of pyridine to
the one-electron-oxidized complexes 1+ and 2 in CH3OH
does not induce significant changes in their visible spectra,
thus confirming that these species are already octahedral
NiIII complexes with two solvent molecules bound axially.
However, 3ꢀ exhibits a quite different behavior: the oxi-
dized product in CH3OH at 233 K exhibits strong absorp-
tions at 821 (e=2600 mꢀ1 cmꢀ1) and 327 nm (e=15400) and
unresolved bands in the region 350–500 nm. These bands,
which match reasonably well with those observed in neat
CH2Cl2 for the radical complex, disappear upon addition of
excess pyridine to give a more typical octahedral NiIII spec-
trum. 3ꢀ thus has a radical character both in CH2Cl2 and
CH3OH. Conversion of this radical species into the corre-
sponding NiIII complex requires a strong N-donor exogenous
ligand (CH3OH is known to have only a moderate coordi-
nating ability).
The EPR spectra of the one-electron-oxidized complexes
1+, 2, and 3ꢀ in CH3OH at 100 K, in the absence of pyridine,
show rhombic low-spin NiIII signals. The geometry around
the metal center is expected to be similar to that in the mix-
ture of CH2Cl2 and pyridine since the pattern of g values is
similar. As expected, no hyperfine coupling could be detect-
ed in any of the g component (the 16O isotope has no nucle-
ar spin). The gx and gy values obtained for the CH3OH
adduct are higher than those measured for the pyridine
adduct, which reflects a stronger interaction between the
metal dz2 orbital and the axial pyridine lone-pair that desta-
bilizes the nickel SOMO.[23]
It has recently been shown[9] that the square-planar nick-
el(ii) complex [NiII(MePy(Bz)2)(tBu2Cat)]ꢀ (MePy(Bz)2 =
N,N-bis(benzyl)-N-[(6-methyl-2-pyridyl)methyl]amine;
tBu2Cat=3,5-di-tert-butylcatecholate) can be oxidized in a
one-electron process, in CH2Cl2, into a nickel(iii) catecholate
species (7), which is converted into the nickel(ii) radical
complex 8 (Scheme 2) in the presence of exogenous ligands.
This behavior contrasts sharply with that observed for 1+, 2,
and 3ꢀ where, upon addition of strong exogenous ligands,
the delocalized radical species are converted into nickel(iii)
phenolate complexes. In these compounds, the nickel(iii) ion
resides in an octahedral geometry, with two pyridines bound
axially (Scheme 2).
These results illustrate the high level of control of the oxi-
dation locus that can be reached simply by modulating the
ligand field. In addition, we have shown how it tunes the
stability of the radical species, as well as its affinity for exog-
enous ligands.
NiIII species are thus systematically observed in the mix-
ture of CH2Cl2 and pyridine, while delocalized radicals are
observed in CH2Cl2 alone. Axial ligation of pyridine shifts
the dz2 orbital above the filled ligand orbitals and promotes
an electron transfer from the metal to the open-shell ligand.
Both 1+ and 2 have a NiIII character in the coordinating
solvent CH3OH, which can be explained by axial ligation of
two solvent molecules in a similar manner to pyridine. In
contrast, 3ꢀ exhibits an EPR spectrum corresponding to an
ExperimentalSection
General: All chemicals were of reagent grade and were used without pu-
rification. Microanalyses were performed by the Service Central d’Anal-
yses du CNRS (Lyon, France).
NMR spectra were recorded with a Bruker AM 300 spectrometer (1H at
300 MHz, 13C at 75 MHz). Chemical shifts are given relative to tetrame-
thylsilane (TMS). Mass spectra were recorded with a Thermofunnigen
(EI/DCI) apparatus.
Chem. Eur. J. 2006, 12, 2293 – 2302
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2299

F. Thomas, O. Jarjayes et al.
butyl-2-hydroxybenzaldehyde
(two
equivalents) in ethanol according to
the procedure published by Wçltinger
et al.[25]
3,5-Di-tert-butylsalicyl chloride: 3,5-
Di-tert-butylsalicylic
acid
hydrate
(1.0 g, 4 mmol) was dissolved in THF
(50 mL). Thionyl chloride (5.9 mL,
80 mmol) was slowly added under ni-
trogen, while cooling in an ice bath,
and the mixture was stirred overnight
and then allowed to warm to room
temperature. The solvent and the
excess thionyl chloride were evaporat-
ed in vacuo and the resulting product,
isolated as a pale-yellow oil, was used
directly in the next step without fur-
ther purification. 1H NMR (300 MHz,
CDCl3, 298 K, TMS): d=10.24 (s,
4
1H), 7.91 (d, J=2.3 Hz, 1H), 7.66 (d,
4J=2.3 Hz, 1H), 1.42 (s, 9H),
1.33 ppm (s, 9H).
Scheme 2. Modulation of the oxidation locus by coordination of an exogenous ligand (Ex) to 3ꢀ (this work)
and 7 (reference [9]).
2-Amino-1-(3,5-di-tert-butyl-2-hydrox-
ybenzimino)benzene: This precursor
for H3L2 was prepared according to
the procedure of Atwood et al.[26]
UV/Vis spectroscopy: UV/Vis spectra were recorded at 298 K with a
Perkin–Elmer Lambda 2 spectrophotometer equipped with a tempera-
ture controller unit set at 298 K. The path length of the quartz cell was
1.000 cm. Low-temperature visible spectra were recorded with a Cary 50
spectrophotometer equipped with a Hellma low-temperature immersion
probe (1.000-cm path-length quartz cell).
2-(3,5-Di-tert-butyl-2-hydroxybenzamido)-1-(3,5-di-tert-butyl-2-hydroxy-
benzimino)benzene (H3L2): As described below for the synthesis of
ligand H4L3, 3,5-di-tert-butylsalicyl chloride (410 mg, 1.5 mmol) in di-
chloromethane (20 mL) was coupled with 2-amino-1-(3,5-di-tert-butyl-2-
hydroxybenzimino)benzene (500 mg, 1.5 mmol) in dichloromethane
(30 mL) and worked up in the same way. The raw product was purified
by column chromatography over silica gel (100 g) with pentane/dichloro-
methane (1:1) to yield H3L2 as a beige powder (340 mg, 41%). 1H NMR
(300 MHz, CDCl3, 298 K, TMS): d=12.43 (br. s, 1H), 12.27 (s, 1H), 8.57
(s, 1H), 8.51 (s, 1H), 8.31 (dd, 3J=8.1, 4J=1.1 Hz, 1H), 7.43–7.41 (m,
2H), 7.3–7.0 (m, 5H), 1.36 (s, 9H), 1.35 (s, 9H), 1.25 (s, 9H), 1.19 ppm (s,
9H); UV/Vis (CH2Cl2): lmax (e)=328 nm (18980 mꢀ1 cmꢀ1); IR (neat): n˜ =
3476 w (nOH), 3427 (nNH), 2960, 2904, 2866 w (nCH), 1646 (nC=O), 1616,
1586 (nC=N) cmꢀ1; MS: m/z (%): 557 (100) [M+H]+.
1,2-Bis(3,5-di-tert-butyl-2-hydroxybenzamido)benzene (H4L3): o-Phenyl-
enediamine (220 mg, 2 mmol) was dissolved in dry dichloromethane
(20 mL) and the solution was cooled to ꢀ408C. A solution of freshly pre-
pared 3,5-di-tert-butylsalicyl chloride (ca. 4 mmol) in dry dichlorome-
thane (50 mL) was then added dropwise to the cooled solution of the di-
amine, under nitrogen, then the cooling bath was removed and the mix-
ture was stirred for a further two hours while it was allowed to warm to
room temperature. The mixture was then filtered to remove a small
amount of a white precipitate, washed with water (2”20 mL) and saturat-
ed NaCl solution (20 mL), and dried with sodium sulfate. The solvent
was evaporated and the raw product recrystallized from methanol to
yield H4L3 as a white powder (0.45 g, 40%). 1H NMR (300 MHz, CDCl3,
298 K, TMS): d=7.56 (dd, 3J=6.0, 4J=3.5 Hz, 2H), 7.51 (d, 4J=2.2 Hz,
2H), 7.38 (d, 4J=2.2 Hz, 2H), 7.30 (dd, 3J=6.0, 4J=3.5 Hz, 2H), 1,43 (s,
18H), 1,28 ppm (s, 18H); UV/Vis (CH2Cl2): lmax (e)=324 nm
(12320 mꢀ1 cmꢀ1); IR (neat): n˜ =3428 w (nOH), 3303 (nNH), 2957, 2908,
2870 w (nCH), 1624 (nC=O) cmꢀ1; MS: m/z (%): 573 (100) [M+H]+.
The affinity constants were obtained by refinement of the UV/Vis titra-
tion data of the complexes with pyridine in CH2Cl2. The fit was per-
formed with the commercial SPECFIT software according to Equation (1).
EPR spectroscopy: X-band EPR spectra were recorded with a Bruker
ESP 300E spectrometer equipped with a Bruker nitrogen-flow cryostat.
Spectra were treated by using the WINEPR software and simulated
using the Bruker SIMFONIA software. Spectra were recorded at 100 K
with 200-mL samples, while variable-temperature experiments were per-
formed with 50-mL samples.
Electrochemistry: The cyclic and differential pulse voltammograms of
each compound (1 mm) in CH2Cl2 containing 0.1m tetra-n-butylammoni-
um perchlorate (TBAP) as supporting electrolyte were recorded with a
CHI 660 potentiostat at 298 K using a glassy carbon disc as working elec-
trode, a Pt wire as secondary electrode, and an Ag/AgNO3 (0.01m) refer-
ence electrode. The potential of the ferrocenium/ferrocene (Fc+/Fc)
couple was used as an internal reference (+0.087 V under our experi-
mental conditions). Experiments were performed under argon. Electroly-
sis at a carbon-felt electrode was performed at 233 K, by controlled po-
tential electrolysis, with a PAR 273 potentiostat. It was followed by rotat-
ing disc voltammetry (carbon-disc electrode, 600 rpm).
Crystalstructure anaylsis : For all structures, the collected reflections
were corrected for Lorentz and polarization effects but not for absorp-
tion. The structures were solved by direct methods and refined by using
the teXsan software package (texsan, Crystal Structure Analysis Package,
Molecular Structure, Corp., The Woodlands, TX, 1992). All non-hydro-
gen atoms were refined with anisotropic thermal parameters. Hydrogen
atoms were generated in idealized positions, riding on the carrier atoms,
with isotropic thermal parameters. CCDC-269657 ([Ni(L1)]), CCDC-
269762 ([Ni(L2)]ꢀ), and CCDC-269450 ( [Ni(L3)]2ꢀ) contain the supple-
mentary crystallographic data for this paper. These data can be obtained
free of charge from the Cambridge Crystallographic Data Centre via
www.ccdc.cam.ac.uk/data_request/cif.
[Ni(L1)]: H2L1 (50 mg, 90 mmol) was dissolved in ethanol (40 mL) and a
solution of Ni(OAc)2·4H2O (45 mg, 180 mmol) in ethanol (30 mL) was
added, followed by triethylamine (26 mL, 180 mmol). The mixture was re-
fluxed for three hours, then two thirds of the solvent was evaporated and
the solution stored at ꢀ208C overnight. [Ni(L1)] (44 mg, 80%) was fil-
tered off as fine, red needles and washed with ice-cold ethanol. Elemen-
tal analysis calcd (%) for C36H46N2NiO2: C 72.37, H 7.76, N 4.69, Ni 9.82;
found: C 71.98, H 7.72, N 4.61, Ni 9.50; 1H NMR (300 MHz, CD2Cl2,
3
4
Preparation of the ligands and complexes
1,2-Bis(3,5-di-tert-butyl-2-hydroxybenzimino)benzene (H2L1): H2L1 was
prepared by refluxing o-phenylenediamine (one equivalent) and 3,5-tert-
298 K, TMS): d=8.33 (s, 2H), 7.78 (dd, J=6.3, J=3.3 Hz, 2H), 7.46 (d,
4J=2.6 Hz, 2H), 7.27 (dd, 3J=6.3, 4J=3.3 Hz, 2H), 7.20 (d, 4J=2.6 Hz,
2H), 1,49 (s, 18H), 1,36 ppm (s, 18H); UV/Vis (CH2Cl2): lmax (e)=
2300
www.chemeurj.org
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Chem. Eur. J. 2006, 12, 2293 – 2302

Nickel Complexes of Pro-Radical Ligands
FULL PAPER
383 nm (29600 mꢀ1 cmꢀ1), 492 (9200); IR (neat): n˜ =2959, 2897, 2864 w
(nCH), and 1616, 1605, 1583 s (nC=N) cmꢀ1. MS: m/z (%): 597 (100)
[M+H]+.
Chem. 1991, 30, 4654–4663; J. Seth, V. Palaniappan, D. F. Bocian,
Inorg. Chem. 1995, 34, 2201–2206; M. W. Renner, J. Fajer, J. Biol.
Inorg. Chem. 2001, 6, 823–830.
[Ni(L2)](Bu4N): Ni(OAc)2·4H2O (45 mg, 180 mmol) in ethanol (30 mL)
was added to a solution of H3L2 (50 mg, 90 mmol) in ethanol (20 mL).
After further addition of triethylamine (50 mL, 360 mmol) and Bu4NOH
(90 mL of a 1m solution in methanol, 90 mmol), the mixture was refluxed
for four hours. Two thirds of the solvent was then evaporated and, after
storage overnight at ꢀ208C, [Ni(L2)](Bu4N) was obtained as a dark-
orange precipitate. Single crystals were obtained by slow evaporation of
the ethanol (51 mg, 76%). Elemental analysis calcd (%) for
C52H81N3NiO3·H2O: C 71.55, H 9.58, N 4.81; found: C 71.17, H 9.53, N
4.85; 1H NMR (300 MHz, CD3OD, 298 K, TMS): d=8.7–8.65 (m, 1H),
8.33 (s, 1H), 7.80 (d, 4J=2.6 Hz, 1H), 7.7–7.6 (m, 1H), 7.21 (d, 4J=
2.5 Hz, 1H), 7.11 (d, 4J=2.5 Hz, 1H), 7.06 (d, 4J=2.6 Hz, 1H), 6.95–6.87
(m, 1H), 6.78–6.71 (m, 1H), 3.21–3.07 (m, 8H), 1.60–1.43 (m, 8H), 1.35
(s, 9H), 1.34 (s, 9H), 1.32–1.22 (m, 8H), 1,20 (s, 9H), 1.19 (s, 9H),
0.87 ppm (t, 3J=7.3 Hz, 12H); UV/Vis (CH2Cl2): lmax (e)=325 nm
(10600 mꢀ1 cmꢀ1), 366 (10400), 427 (9400); IR (neat): n˜ =2955, 2896, 2864
w (nCH), 1594 s (nC=O), and 1561, 1521 (nC=N) cmꢀ1. MS: m/z (%): 611
(100) [Ni(L2)]ꢀ.
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[13] [Zn(L1)] is the zinc(ii) complex of H2L1 obtained by mixing Zn-
(ClO4)2·6H2O and the ligand in ethanol in the presence of NEt3.
The phenoxylzinc radical species [Zn(L1)]+ exhibits an unresolved
isotropic (S=1/2) signal at giso =2.005 at 260 K. [Zn(L1)]2+ is a di-
radical species where the two spins are antiferromagnetically ex-
change-coupled (J estimated to be about ꢀ13 cmꢀ1); O. Rotthaus, O.
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[Ni(L3)](Bu4N)2: nBu4NOH (360 mL of
a 1m solution in methanol,
360 mmol) was added to a solution of H4L3 (50 mg, 87 mmol) and Ni-
(OAc)2·4H2O (26 mg, 104 mmol) in DMF (5 mL), and the reaction mix-
ture was stirred at 808C for four hours. The DMF was then distilled off
in vacuo and the raw product was purified over LH 20 (15 g) with metha-
nol as eluent. After evaporation of the solvent, [Ni(L3)](Bu4N)2 was ob-
tained as a red powder (47 mg, 48%). Single crystals were obtained by
slow evaporation of a diethyl ether/pentane solution. Elemental analysis
calcd (%) for C68H116N4NiO4·2H2O: C 71.12, H 10.53, N 4.88, Ni 5.11;
found: C 72.08, H 10.53, N 4.88, Ni 4.97. 1H NMR (300 MHz, CD2Cl2,
298 K, TMS): d=8.48 (dd, 3J=6.2, 4J=3.6 Hz, 2H), 7.77 (d, 4J=2.8 Hz,
2H), 6.69 (d, 4J=2.8 Hz, 2H), 6.43 (dd, 3J=6.2, 4J=3.6 Hz, 2H), 3.25–
3.19 (m, 16H), 1.42–1.21 (m, 34H), 1,18 (s, 18H), 1.06–1.00 (m, 16H),
0.68 ppm (t, 3J=7.2 Hz, 24H). UV/Vis (CH2Cl2): lmax (e)=341 nm
(13300 mꢀ1 cmꢀ1), 368 (14500), 391 sh (11400) 515 (600). IR (neat): n˜ =
2951, 2871 w (nCH) and 1604 s (nC=O) cmꢀ1; MS: m/z (%): 869 (100)
[{Ni(L3)}(Bu4N)]ꢀ, 627 (51) [{Ni(L3)}H]ꢀ, 313 [Ni(L3)]2ꢀ
.
[15] A similar experiment performed in CH3NO2 revealed a behavior
that is quite similar to that in CH2Cl2. The TC is, however, shifted to
240 K, a value corresponding to the melting point of the mixture
CH3NO2 + 0.1m TBAP.
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[17] This may be interpreted in different ways, such as fast relaxation,
changes in electronic distribution, or dimerization. We tend to favor
the former hypothesis since a similar behavior (EPR silence at
298 K) is observed for the pyridine adducts [Ni(L2)(Py)2] and
[Ni(L3)(Py)2]ꢀ.
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(mꢀ2) for Equation (1) as a function of T are: for 1+: 6.71ꢂ0.09
(298 K), 7.94ꢂ0.14 (263 K), and 8.76ꢂ0.18 (243 K); for 2: 7.26ꢂ
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ꢀ 2006 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemeurj.org
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F. Thomas, O. Jarjayes et al.
0.13 (233 K); for 3ꢀ: 2.41ꢂ0.06 (298 K), 4.04ꢂ0.09 (263 K), and
5.43ꢂ0.10 (238 K). The parameters DH° and DS° could be obtained
from these values: DH° =51.7 kJmolꢀ1 and DS° =ꢀ44.9 JmolKꢀ1
for 1+ and DH° =68.2 kJmolꢀ1 and DS° =ꢀ183 JmolKꢀ1 for 3ꢀ.
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[22] Owing to the confidence level of ꢂ0.05 mT on Ax and Ay, the un-
paired spin density on the pyridine 14N could not be calculated accu-
rately. In the 1:1 [Ni(L1)]+:pyridine mixture we were able to trap
the 1:1 pyridine adduct, which is characterized by a triplet (instead
of a quintuplet) in the low-field component. As expected, the Azz
value (2.5 mT) of [Ni(L1)(Py)]+ is higher than that of [Ni(L1)(Py)2]+.
From the UV/Vis titration, a logb value of 4.15ꢂ0.05 was obtained
at 298 K for the equilibrium [Ni(L1)]+ +PyQ[Ni(L1)(Py)]+.
I. C. Santos, M. Vilas-Boas, M. F. M. Piedade, C. Freire, M. T.
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chemically (the cations 2+ and 3+ were too unstable to be generated
electrochemically). However, the low stability of their pyridine ad-
ducts, and the lack of X-band EPR signals, makes the assignment of
the oxidation locus difficult (see Supporting Information).
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Received: July 29, 2005
Revised: September 28, 2005
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Published online: December 21, 2005
2302
www.chemeurj.org
ꢀ 2006 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Chem. Eur. J. 2006, 12, 2293 – 2302