Spectroscopic and electrochemical properties of nickel(II), iron(III) and cobalt(II) complexes with benzilbisthiosemicarbazone - Importance of working conditions and the metal salt used in the final complex

Article abstract of DOI:10.1016/S0277-5387(00)00505-2

Reactions of benzilbisthiosemicarbazone (LH₆) with nickel, cobalt and iron chloride and nitrate give different complexes depending on the salts used and the working conditions. Reactions from nickel chloride give three complexes, 1, 2 and 3 when the solvent or the pH is modified. Complex 1 is obtained in ethanol under reflux, and in methanol, independently of temperature and also with basic medium. Complex 2 is formed by working in ethanol at room temperature. The ligand acts as a dianion in both complexes. In the presence of hydrochloric acid, the new isolated complex 3 presents the neutral form of the ligand. Reaction from nitrate yields only complex 1. Reactions with iron chloride give two new complexes with different formulae and grade of deprotonation in the ligand. For complex 4, obtained at room temperature, LH₆ acts as an anion with a ligand-iron ratio of 2:1 and in its neutral form with a 1:1 ratio in complex 5, working under reflux. Reactions with iron nitrate yield complex 6, with the ligand in an intermediate situation, as anion but a nitrate remains. Complexes 7 and 8, with a 1:1 cobalt-ligand molar ratio are isolated from reactions with cobalt nitrate; both contain nitrate and ethanol or water to complete the co-ordination sphere. Attempts to get a macrocyclic thiosemicarbazone from reaction of benzil and LH₆ in the presence of nickel or iron salt gave the complexes obtained, in absence of the dicarbonyl molecule. Electrochemical behaviour of complexes studied by cyclic voltammetry show metal-centred reduction processes for all of them. The reduction/oxidation potential values depend on the structures of complexes. Nickel complexes exhibit waves corresponding to Ni(III)-Ni(II) process. Electrochemical response of iron complexes depends on the presence of chloride ions. (C) 2000 Elsevier Science Ltd.

Full text of DOI:10.1016/S0277-5387(00)00505-2

www.elsevier.nl/locate/poly
Polyhedron 19 (2000) 2059–2068
Spectroscopic and electrochemical properties of nickel(II), iron(III)
and cobalt(II) complexes with benzilbisthiosemicarbazone —
importance of working conditions and the metal salt used in the
final complex
Marta Can˜adas a, Elena Lo´pez-Torres a, Arturo Mart´ınez-Arias b,
Ma. Antonia Mendiola a,*, Ma. Teresa Sevilla c
a Departamento de Qu´ımica Inorga´nica, Uni6ersidad Auto´noma de Madrid, 28049 Madrid, Spain
b Instituto de Cata´lisis y Petroleoqu´ımica, CSIC, Campus UAM, 28049 Madrid, Spain
c Departamento de Qu´ımica Anal´ıtica y Ana´lisis Instrumental, Uni6ersidad Auto´noma de Madrid, 28049 Madrid, Spain
Received 17 February 2000; accepted 10 July 2000
Abstract
Reactions of benzilbisthiosemicarbazone (LH6) with nickel, cobalt and iron chloride and nitrate give different complexes
depending on the salts used and the working conditions. Reactions from nickel chloride give three complexes, 1, 2 and 3 when
the solvent or the pH is modified. Complex 1 is obtained in ethanol under reflux, and in methanol, independently of temperature
and also with basic medium. Complex 2 is formed by working in ethanol at room temperature. The ligand acts as a dianion in
both complexes. In the presence of hydrochloric acid, the new isolated complex 3 presents the neutral form of the ligand. Reaction
from nitrate yields only complex 1. Reactions with iron chloride give two new complexes with different formulae and grade of
deprotonation in the ligand. For complex 4, obtained at room temperature, LH6 acts as an anion with a ligand–iron ratio of 2:1
and in its neutral form with a 1:1 ratio in complex 5, working under reflux. Reactions with iron nitrate yield complex 6, with the
ligand in an intermediate situation, as anion but a nitrate remains. Complexes 7 and 8, with a 1:1 cobalt–ligand molar ratio are
isolated from reactions with cobalt nitrate; both contain nitrate and ethanol or water to complete the co-ordination sphere.
Attempts to get a macrocyclic thiosemicarbazone from reaction of benzil and LH6 in the presence of nickel or iron salt gave the
complexes obtained, in absence of the dicarbonyl molecule. Electrochemical behaviour of complexes studied by cyclic voltammetry
show metal-centred reduction processes for all of them. The reduction/oxidation potential values depend on the structures of
complexes. Nickel complexes exhibit waves corresponding to Ni(III)–Ni(II) process. Electrochemical response of iron complexes
depends on the presence of chloride ions. © 2000 Elsevier Science Ltd. All rights reserved.
Keywords: Thiosemicarbazones; Electrochemistry; Iron complexes; EPR of iron complexes; Nickel complexes; Cobalt complexes
with many of the first row transition metals have shown
a wide range of biological properties [3–5].
1. Introduction
It has been shown that a NNS tridentate system is
present in most of the thiosemicarbazones with carcino-
static potency [6] and that copper and iron complexes
usually have higher activity than the free ligands [7,8].
The chemical properties of these complexes, namely
concerning their behaviour in redox reactions, relate to
those observed in biological systems. For some copper
complexes, a lower reduction potential seems to be
related to an increased antifungal activity and with
their potential abilities as SOD–mimic complexes [8–
The investigation of metal complexes with sulfur-
containing Schiff bases is a subject of current and
growing interest and it has been shown that many of
them present anticancer activity [1,2]. A particular im-
portant group of such ligands includes those containing
thiosemicarbazone moieties, because complexes formed
* Corresponding author. Tel.: +34-91-397-4844; fax: +34-91-397-
4833.
E-mail address: antonia.mendiola@uam.es (M. Antonia Mendiola).
0277-5387/00/$ - see front matter © 2000 Elsevier Science Ltd. All rights reserved.
PII: S0277-5387(00)00505-2

2060
M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
2.2. Synthesis of nickel complexes
2.2.1. NiC16H14N6S2 (1)
A solution of nickel (II) chloride (0.20 g, 0.8 mmol)
in methanol (10 cm3) was added to a suspension of LH6
(0.30 g, 0.8 mmol) in 20 cm3 of the same solvent. The
mixture was stirred for 3 h at room temperature. The
reddish solid formed was filtered off, washed with
methanol and dried in vacuo (yield 73%). 13C NMR
(DMSO-d6): l(ppm) 127.7, 128.8, 129.4, 131.6 (CPh),
153.8 (CN), 179.7 (CS). H NMR (DMSO-d6): l(ppm)
7.6 (s, NH2), 7.1–7.3 (m, HPh).
The same solid was obtained in ethanol under reflux
[27]; in the presence of LiOH (0.028 g, 0.8 mmol) under
reflux; and from nickel nitrate in ethanol at room
temperature.
Fig. 1. Proposed structure for benzilbisthiosemicarbazone, LH6.
11]. Recently, iron-based SOD mimics have been re-
ported [12–14]. The most successful mimics were ob-
tained using macrocyclic pentadentate ligands [12].
However, these structures are quite different from the
distorted trigonal bipyramidal geometry of the active
site structure of Fe–SODs [15]. Some iron complexes
with tripodal tetradentate ligand have activity [16]. The
mimetic activity is related to the reduction potential
values in the Fe(III)–Fe(II) process. The values ob-
served in most Fe–SOD are between 0.2 and 0.4 V
(versus NHE) [17]. The factors that influence the acces-
sibility of several oxidation states for nickel are cur-
rently studied, because it is related to the bio-inorganic
role of this metal ion in several redox enzymes [18,19].
Further interest in the co-ordination chemistry of nickel
(I) and cobalt (I) arise from the role of these complexes
in several catalytic reactions, such as electrocatalytical
[20–22] and photocatalytical [23,24] reduction of CO2.
The incorporation of both soft and hard donor atoms
in the co-ordination environment of the metal ion may
be crucial in the stabilisation of oxidation numbers +1
and +3 for nickel [25].
Several studies on metal complexes of ligands derived
from benzil and thiohydrazides with cobalt (II), zinc
(II) and nickel (II) chloride and copper (II) salts have
been made in our group [26,27]. Copper complexes
have also been studied as SOD mimetic types [28,29]. In
this paper we report the details of the synthesis, the
spectroscopic characterization and the electrochemical
behaviour of complexes obtained from benzil-
bisthiosemicarbazone (Fig. 1) and nickel (II), cobalt (II)
and iron (III) salts. We have also investigated the
influence of the counter-ion and the pH of the reaction
medium in the behaviour of the ligand in reactions with
metal chloride or nitrate.
1
2.2.2. NiC32H30N12S4 (2)
This complex was prepared in the same way using
ethanol and at room temperature (yield 63%) [27].
2.2.3. NiC16H16N6S2Cl2 (3)
This complex was obtained following the same proce-
dure, but in the presence of 10 cm3 of hydrochloric acid
(12 M) in methanol and under reflux (yield 58%).
2.2.4. Reactions with benzil
Working in the conditions described, but in the pres-
ence of benzil (0.17 g, 0.8 mmol) in methanol. The
mixture was (a) refluxed with stirring for 3 h; (b)
refluxed for 3 h in the presence of lithium hydroxide
(0.034 g, 0.8 mmol); and (c) refluxed for 3 h in the
presence of 10 cm3 of hydrochloric acid (12 M). A
reddish brown solid 1 was formed in procedures (a) and
(b) but a brown solid 3 was formed following (c).
2.3. Synthesis of iron (III) complexes
2.3.1. FeC32H29N12S4 (4)
A solution of iron (III) chloride (0.23 g, 0.8 mmol) in
10 cm3 of methanol was added to a suspension of LH6
(0.30 g, 0.8 mmol) in 20 cm3 of the same solvent. The
mixture was stirred for 6 h at room temperature. The
solid formed was filtered off, washed with methanol
and dried in vacuo (yield 31%).
2. Experimental
2.3.2. FeC17H20N6S2OCl3 (5)
2.1. Reagents
The reaction took place in the same conditions, but
heating the mixture under reflux for 6 h gave a green
solution, which was concentrated until a dark brown
solid was formed (yield 43%). A reaction in the pres-
ence of hydrochloric acid in the reaction medium,
enabled the reagents to be isolated. The reaction in the
presence of LiOH (0.034 g, 0.8 mmol) in methanol
under reflux for 2 h gave a yellow solid (yield 21%) [30].
Nickel(II) and cobalt(II) chloride and nitrate hexahy-
drate, iron(III) chloride hexahydrate and iron (III) ni-
trate nonahydrate were commercial products of highest
chemical grade (Fluka). Solvents were purified accord-
ing to standard procedures.
Benzilbisthiosemicarbazone LH6 was prepared fol-
lowing the method previously reported [26].

M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
2061
2.3.3. Reactions with benzil
Following the same procedure described for nickel
(II) chloride. Using procedures (a) and (c), the reagents
were recovered, but in the presence of lithium hydrox-
ide the yellow solid obtained above was isolated.
2.3.4. FeC16H20N7S2O6 (6)
A solution of iron (III) nitrate (0.52 g, 1.3 mmol) in
ethanol (10 cm3) was added to a suspension of LH6
(0.46 g, 1.3 mmol) in 40 cm3 of ethanol. Immediately
the colour of the solution changed to dark brown, then
the mixture was stirred for 8 h. The dark solid was
filtered off, washed with ethanol and dried in vacuo
(yield 27%).
2.4. Synthesis of cobalt(II) complexes
2.4.1. CoC18H21N7S2O4 (7)
A solution of cobalt(II) nitrate (0.54 g, 1.8 mmol) in
ethanol (25 cm3) was added to a suspension of LH6
(0.66 g, 1.8 mmol) in 40 cm3 of ethanol. Immediately
the solution changed to a dark brown colour. The
mixture was stirred for 8 h. The dark solid was filtered
off, washed with ethanol and dried in vacuo (yield
66%).
2.4.2. CoC16H21N7S2O6 (8)
The same procedure was followed, as above, but
heating the mixture under reflux with stirring for 8 h
(yield 82%).
2.5. Physical measurements
Microanalyses were carried out using a Perkin–
Elmer 2400 II CHNS/O elemental analyser. IR spectra
in the 4000–400 cm−1 range were recorded as KBr
pellets on a Bomen-100 spectrophotometer and in the
range 550–200 cm−1 on a Bruker IFS 66V.
UV–Vis spectra in solution were registered on an
Ati-Unicam UV2 spectrophotometer and the solid
reflectance spectra on a Pye-Unicam SP-8-100. Conduc-
tivity data were measured using freshly prepared DMF
solutions (ca. 10−3M) at 25°C with a Metrohm Herisau
model E-518 instrument.
1H and 13C NMR spectra were recorded on a Bruker
AMX-300 spectrophotometer, using DMSO-d6 as a
solvent and Me4Si as an internal standard.
Fast atom bombardment mass spectra were recorded
on a VG Auto Spec instrument using Cs as the fast
atom and m-nitrobenzylalcohol (mNBA) as the matrix.
EPR spectra were obtained with a Bruker ER 200D
spectrometer operating in the X-band and calibrated
with a DPPH standard (g=2.0036). Portions of the
samples were placed inside a quartz probe cell and
subjected to brief out-gassing at room temperature (in
high vacuum conditions) prior to record the spectra.

2062
M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
Table 2
Relevant IR spectral data and assignament of LH6 and its complexes (cm−1
)
wNH+wNH2
wCꢀN+lNH2
Thioamide I
1465 a
Thioamide IV
848
w(MꢁN)
w(MꢁX)
LH6
3420, 3384 3342, 3328 3236, 1613, 1585
3145
NiLH4 (1)
Ni(LH5)2 (2)
NiLH6Cl2 (3)
3450, 3273 3119
3456, 3419 3279, 3235 3131
3398, 3387, 3273, 3098
1625, 1615, 1610
1444
1474, 1440
1460, 1444
852
869
859
489, 432
506, 432
487, 432
1610, 1602 1588
1615, 1599, 1583,
1570
313,324
382
Fe(LH5)(LH4) (4)
3419, 3260, 3249, 3144
3300–3160
1590, 1581
1479, 1470,
1440
1444
866
849
503, 437, 426
506, 475
FeLH6Cl3 CH3OH
1617, 1507, 1575
(5)
FeLH4NO3·3H2O (6) 3420, 3286 3219 b
1618 b, 1574
1615
1439
1463
844
848
485
481
295 a
340
CoLH5NO3
CH3CH2OH (7)
CoLH5NO3·3H2O
(8)
3300–2700 b
3300–2700 b
1626 b
1482
835
481
342
a Several bands.
b Solvent molecules (water, methanol or ethanol).
Electrochemical measurements were performed with
a BAS CV 27 voltammograph and a BAS A-4 XY
register, using a carbon glassy (ƒ 5 mm) as working
electrode, a platinum wire as auxiliary and a double
junction, with porous ceramic wick, Ag–AgCl as refer-
ence electrode. In all cases, ferrocene was used as an
external standard under the experimental conditions
chloride, under the working conditions described in the
experimental section, were unsuccessful. The complexes
isolated from the nickel salt and benzilbisthiosemicar-
bazone were the same in basic or acid medium.
The nature of iron complexes is very dependent on
the working conditions: temperature, salt used or the
reaction medium. From the iron chloride reactions in
methanol, two complexes 4 (Fe(LH5)(LH4)) and 5
(FeLH6Cl3CH3OH) are obtained depending on the tem-
perature. Complex 4 presents a 1:2 iron–ligand molar
ratio with LH6 acting as an anion, but complex 5 has
chloride ions and the ligand is in its neutral form. In
comparison with the nickel reaction, no complex is
obtained by working in the presence of hydrochloric
acid, while in the basic medium an iron macrocyclic
complex was formed [30]. Moreover, from the nitrate
reaction a new complex 6 (FeLH4NO3·3H2O), which
has a nitrate ion, is isolated.
As in the case of nickel, attempts to get the macro-
cyclic ligand from the reaction of benzil and LH6 in the
presence of iron chloride, under the working conditions
described in the experimental section, were unsuccess-
ful. A macrocyclic complex was obtained working in
the presence of lithium hydroxide. This complex was
obtained in the absence of benzil as shown in a previ-
ous study [30].
employed, redox couple ferricinium–ferrocene (E1/2
=
+0.48 V; DEp=60 mV). Cyclic voltammetry studies of
ligands and complexes were carried out in 0.01 M
solutions in dimethylformamide containing 0.1 M
[NBu4] [PF6] and lithium chloride as supporting elec-
trolytes. The range of potential studied was from +1.5
to −2.5 V. All solutions were purged with nitrogen
steam for 5 min before measurements and the working
electrode was polished before each experiment with
diamond paste. The procedure was performed at room
temperature and nitrogen atmosphere was maintained
over the solution during the measurements.
3. Results and discussion
Analytical data are summarised in Table 1. From
these data it can be deduced that the formulae of
complexes depends on the working conditions.
For nickel reactions, three complexes were obtained:
1 (NiLH4) and 2 (Ni(LH5)2), in which the ligand acts as
an anion and complex 3 (NiLH6Cl2) with LH6 in its
neutral form. The last one is only obtained in the
presence of hydrochloric acid, while complex 1 appears
independently of the salts used, the solvent (methanol
or ethanol), under reflux and in the presence of a basic
medium.
From cobalt nitrate reactions, the obtained com-
plexes 7 (CoLH5NO3CH3CH2OH) and 8 (CoLH5NO3·
3H2O) present the ligand as partially deprotonated as in
complex 6, in contrast to the complex obtained from
chloride [27]. In both cobalt complexes, solvent
molecules are bonded to cobalt, ethanol in complex 7
and water in complex 8.
The conductivity data (Table 1) indicate that all
complexes are non-electrolytes, although complexes 5, 7
and 8 show a partial solvolysis [31].
Attempts to get the macrocyclic ligand from the
reaction of LH6 and benzil in the presence of nickel

M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
2063
3.1. Nickel complexes
those of complex 1, so the behaviour of the ligand must
be the same. The bands present in the far infrared
spectrum of complex 3 support that the imine and
thioamide nitrogen atoms are bonded to the metal ion.
This spectrum shows two additional bands assignable
to w(NiꢁCl) with respect to complexes 1 and 2, which
confirm the chlorine–nickel bonds (Table 2). Electronic
spectrum in DMF solution shows absorption bands at
707 nm (m=333 mol−1 dm3 cm−1) and 744 nm (m=
304 mol−1 dm3 cm−1) and three more intense bands at
430, 416 nm (m=3103 mol−1 dm3 cm−1) and 312 nm
(m=12.244 mol−1 dm3 cm−1) and in solid state at 540
and 680 nm with a shoulder at 740 nm, which are
consistent with a distorted octahedral environment for
the nickel [32].
Mass spectra of all complexes show a peak at 413
amu corresponding to the [LH5Ni]+ fragment. The
spectrum of complex 3 does not exhibit the molecular
ion peak, but since it is a molecular complex the
chloride ions must be bonded to the metal.
1H NMR spectra of complexes 1 and 2 exhibit the
amine and aromatic hydrogen in the same range as the
free ligand. This is diagnostic of the diamagnetism for
complexes, which indicates a square-planar symmetry
for them in DMSO solution.
The spectrum of complex 2 confirms the presence of
two non-equivalent ligand molecules. Its spectrum ex-
hibits signals in the same position as complex 1 and the
free ligand. This data indicates that one of the ligands
is replaced by the DMSO. The 13C NMR spectrum
supports the partial substitution of one of the ligands
by the solvent. The 13C NMR spectrum of 1 shows a
signal corresponding to the thioamide group in the
same position as the free ligand, but the imine group is
shifted to a lower field. These data support both azome-
thine groups and the primary amines being bonded to
the nickel ion, as was proposed from the previous
spectral data [27].
3.2. Iron complexes
The reactions of iron salts with LH6 gave three new
complexes with different stoichiometry and variable
grades of deprotonation in the ligand (Table 1).
From the chloride reaction, complexes 4 and 5 show
both variables. The mass spectrum of 4 presents a
pattern similar to that of complex 2, which had the
same stoichiometry. The IR spectrum is also close, with
poorly resolved bands. Therefore, the same co-ordina-
tion behaviour of the thiosemicarbazone must be ex-
pected in both complexes, in which one ligand acts as
dianion bonding through two imine nitrogen atoms and
the terminal amine groups, and the other by the amine
The IR spectrum of 3 shows several bands between
3398 and 3098 cm−1 corresponding to NꢁH stretching
vibrations, which indicates that the ligand is present in
its neutral form according to the analytical data. The
shifts of the other bands in this spectrum are similar to
Fig. 2. EPR spectra of iron complexes in solid state, recorded at 77 and 298 K.

2064
M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
group [27]. Analytical data for complex 5 agree with a
1:1 molar ratio and the presence of methanol and
chloride ions. Its IR spectrum confirms that the ligand
acts in its neutral form and is bonded by the azome-
thine nitrogen atoms to the iron ion. The bands in the
far IR confirm this behaviour, as do the bonds to the
iron by the methanol molecule and the chloride ions.
Complex 6, which is obtained from iron nitrate,
presents an intermediate situation. The ligand acts as
an anion, but a nitrate ion remains. The mass spectrum
shows a peak corresponding to [LH4Fe]+. The infrared
spectrum confirms the presence of the nitrate group as
a ligand (mono or bidentate) and the water molecules,
so it is more difficult for the assignment of the amine
vibrations. By comparison with the cadmium nitrate
the solid-state crystallites, giving rise to a relatively
large spreading of the g resonant values [34]. In addi-
tion, complex 6 exhibits a signal at gÞ=2.147, and
gꢀꢀ=1.986 which is resolved by a hyperfine splitting of
five lines in the parallel component, with A=29 G,
which shows intensity decreasing from the central line
to the outermost components of the pattern (the inset
of Fig. 2 shows details of this signal obtained in a
shorter range run). This signal is typical of low-spin
iron(III) complexes in a distorted octahedral environ-
ment with axial symmetry. Analysis of the g tensor
values suggests that the axial ligands are weaker than
the equatorial ones and that the dxy orbital is lower in
energy [35,36]. On the other hand, the hyperfine split-
ting pattern shown by this signal is consistent with a
partial delocalization of the unpaired electron density
into two equivalent nitrogen atoms through a superhy-
perfine interaction. Considering the infrared results, the
nitrogen atoms responsible of this effect might be those
of the imine groups. Besides, in the spectrum recorded
at 298 K, the presence of a relatively broad signal with
g value close to ge is observed. This signal can be
attributed to the low-spin iron(III) system, the broaden-
ing arising from the effects of magnetic exchange inter-
actions within adjacent iron(III) species in the
crystallites or a relatively fast high-low spin configura-
tional equilibrium [36].
A similar broad signal prevails in complexes 4 and 5
indicating the presence of low-spin iron(III) configura-
tions in them. In the former case the presence of
high-spin iron(III) is apparent only in the spectrum
recorded at 298 K (revealed by a peak at g%=5.54),
suggesting a greater relative stability of the low-spin
configuration for complex 4, while complex 5 shows the
coexistence of low- and high-spin states of the iron(III)
signal at low recording temperatures. In the case of
complex 4 a narrower iron(III) signal in the low-spin
state is evidenced by a signal at g1=2.142, g2=2.110
and unresolved g3 in the spectrum recorded at 77 K. A
narrow signal at g=2.01, with relatively low intensity,
is also observed in the spectrum of this complex. This
could probably arise from location of unpaired elec-
trons in the organic ligands of the complex or from the
presence of impurity radicals.
complex in which recent X-ray data confirm
a
monodentate behaviour for this group [33], we propose
that the nitrate is bonded by only one oxygen. The
band assigned to the imine group is shifted to lower
frequency, which indicates that these nitrogen atoms
are involved in the bond to the iron. The iron co-ordi-
nation sphere is completed by four oxygen atoms, one
from the nitrate group and the other from water
molecules (Fig. 3(c)). The bands present in the far IR
confirm the bond FeꢁN(NC) and several FeꢁO bonds
(Table 2).
Electronic spectra in solution show an intense band
at 400 nm (m=103 mol−1 dm3 cm−1) with a shoulder at
766 nm (m=180 mol−1 dm3 cm−1). The reflectance
spectrum of 4 only shows a peak at 320 nm and several
shoulders at lower energy. Spectra of complexes 5 and
6 show bands at 610 and 550 nm with a shoulder at
lower energy in each absorption band, which support a
distorted octahedral geometry for the iron ion. Elec-
tronic spectra in solution show charge transfer bands
that obscure d–d bands [32].
3.2.1. EPR spectra
The EPR spectra of the iron(III) complexes in solid
state, recorded at 77 and 298 K, are shown in Fig. 2. In
all cases an overall intensity decrease is produced when
the spectra are recorded at 298 K related to those
recorded at 77 K.
The spectrum of complex 6 exhibits resonances in the
low field region showing a peak at g%=4.3 and a
relatively broad absorption in the g%=9.0–4.0 range.
This is typical of high-spin iron(III) complexes in rhom-
bically distorted low-symmetry positions [34]. For these
cases the zero-field splitting terms remove the degener-
acy of the spin sextet resulting in three Kramer’s dou-
blets; transitions within these states give rise to
resonances in the low field region of the spectra [34].
Observation of a relatively broad resonance in this
region and a narrower line at g%=4.3 is consistent with
the existence of a certain dispersion in the zero-field
splitting terms. This could be caused by strain effects in
3.3. Cobalt complexes
As it was mentioned in the experimental section and
from the analytical data, two new complexes were
obtained from the nitrate reaction with LH6 (Table 1).
Both complexes differ in the solvated molecules, but the
ligand presents the same grade of deprotonation. The
mass spectra show the peak corresponding to the
[LH5Co]+ ion. IR spectra confirm the presence of the
nitrate ion and the solvent molecules. The band as-
signed to thioamide IV remains in the same position as

M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
2065
easier to oxidize complex 3 than complex 1. This fact
agrees with the presence of two chloride ligands in
complex 3. Chloride is a good p-donor, which will
stabilise the higher oxidation state [41–43].
in the free ligand, which indicates that this group is not
involved in the bonding to the cobalt ion. The presence
of solvent molecules does not permit the assignment of
the NH and imine vibrations, but the bands in the far
IR region indicate that the ligand is bonded by the
imine nitrogen atoms (Table 2) [4].
3.4.2. Iron complexes
The solid state electronic spectra of complexes 7 and
8 show bands at 580 and 500 nm with a shoulder at 690
and 680 nm, respectively, these data are consistent with
a pentacoordinated or a distorted octahedral geometry
for the cobalt ion [32].
Voltammetric studies of complexes 4, 5 and 6 were
carried out in dimethylformamide as solvent and using
tetrabutylammonium hexafluorophosphate (TBAHFP)
An example of the different structures proposed for
complexes is shown in Fig. 3, where the output of the
calculation carried out using the Hyperchem program is
plotted.
3.4. Electrochemical beha6iour
The cyclic voltammogram of the ligand in DMF
shows only two irreversible cathodic waves at −1.44
and −1.75 V, which can be attributed to the reduction
of the groups present in the ligand.
3.4.1. Nickel complexes
Cyclic voltammogram of complex 1 shows the same
pattern in glassy carbon electrode as in platinum [27].
The electrochemical behaviour of complex 3 differs
from that complex 1, in the same conditions, and
depends on the potential range explored. On scanning
from 0.0 to −2.5 V, the cyclic voltammogram (Fig.
4(a)) shows two waves in the cathodic scan at −1.10
and −1.67 V, respectively, corresponding to successive
reductions of the Ni (II) ion [37–39], and in the reverse
scan two anodic peaks at −1.61 and −1.05 V. Analy-
sis of the first redox couple in the interval 25–800 mV
s−1 shows a lineal variation of ipc versus 61/2. Small
dependence of the potential peak with the scan rate is
observed, and the ipa–ipc ratio is close to the unity.
These results suggest a diffusion controlled quasirre-
versible Ni(II)–Ni(I) process. The more negative redox
couple shows similar features, except that the potential
peak values are independent of the scan rate, so this
couple corresponds to a reversible process. When the
positive potential limit is increased to +1.5 V (Fig.
4(b)), complex 3 exhibits an irreversible oxidation at
+0.87 V, which generates a new unknown species
which is turned into an irreversible reduction at −0.67
V [39–41].
Reduction potentials values of complexes 1 and 3 are
similar, although slightly more negative for the latter
complex due to the lower number of unsaturations for
complex 3. It is more difficult to have access to the +1
state for nickel in the complex with the ligand in its
neutral form and two chloride ions bonded to the metal
ion. The values of the oxidation process show that it is
Fig. 3. Possible structures for (a) complex 2; (b) complex 3; (c)
complex 6.

2066
M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
due to the reduction of different species of iron(III)
(Fig. 6(a)).
The three complexes show a different behaviour after
the addition of LiCl. The voltammogram of complex 4
does not change, but those of complexes 5 and 6 are
modified by the presence of chloride ions (LiCl). A
gradual addition of chloride ions yielded a better
defined cyclic voltammogram of complex 5, which has
chloride ions in its formula; a fully reversible Fe(III)–
Fe(II) process is observed, although the potential values
are shifted to +0.04 V. At the level of 0.1 M LiCl a well
defined voltammogram is obtained (Fig. 5(b)). For
complex 6, a dramatic change in the voltammogram is
produced by successive additions of LiCl, suggesting
strong interaction of the complex with the anion by the
replacement of the co-ordinated nitrate ion by chloride.
From the first addition, a clear reversible process at-
tributed to the Fe(III)–Fe(II) couple at +0.03 V and
two peaks at more negative potential are observed (Fig.
6(b)). The features of the voltammogram depend on
chloride content in the solution. A new peak at −0.62
V arises and its intensity increases when the concentra-
tion of chloride increases, so it can be related to the
reduction of mixed species containing iron, ligand and
chloride, [Fe(LH5)Cl3)]−, [Fe(LH5)2)]+, which are re-
sponsible of the electrochemical response observed in
these conditions as previously described [44]. The redox
process Fe(III)–Fe(II) would be accompanied by slow
loss of nitrate, and therefore addition of chloride ions
will improve this and hence more reversible electro-
chemical behaviour is observed. The voltammogram
obtained at the level of 0.1 M LiCl is shown in Fig. 6(c).
Fig. 4. Cyclic voltammogram for complex 3; (a) potential range 0.0 to
−2.5 V; (b) potential range +1.5 to −2.5 V, 6=100 mV s−1
.
and lithium chloride as electrolytes. In addition, voltam-
metric response in TBAHFP was recorded with succes-
sive additions of LiCl until ca. 0.1 M.
In the solutions containing TBAHFP, complexes 4, 5
and 6 show different electrochemical behaviour. The
cyclic voltammogram of complex 4 shows waves at-
tributed to the ligand and poorly defined waves at less
negative potentials. Using 2.10−2 M solutions the cyclic
voltammogram obtained shows two waves in the range
from +1 to −1.7 V. The peak at −0.75 V could be
attributed to a Fe(III)–Fe(II) process. This behaviour
could be explained considering the two ligands around
the metal ion, without flexibility to accommodate the
structural changes that accompany the reduction pro-
cess. Cyclic voltammogram of complex 5 exhibits metal
centred waves and other ones corresponding to the
ligand reduction processes. In the range from +1 to
−1 V a quasirreversible process attributed to the
Fe(III)–Fe(II) couple at +0.19 V and a reduction wave
at −0.72 V are observed (Fig. 5(a)). Cyclic voltammet-
ric response of complex 6 shows poorly defined peaks
Fig. 5. Cyclic voltammogram for complex 5; (a) in TBAHFP; (b) at
the level of 0.1 M LiCl, 6=100 mV s−1
.

M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
2067
Fig. 7. Cyclic voltammogram of complex 7, 6=100 mV s−1
.
related to a quasirreversible process controlled by diffu-
sion. The potential values are slightly shifted with re-
spect to those obtained from the cobalt chloride
complex in the same conditions, because in this com-
plex chloride ions were bonded to the cobalt ion [27].
Fig. 6. Cyclic voltammogram for complex 6; (a) in TBAHFP; (b) in
TBAHFP after the addition of 0.1 ml of LiCl 0.1 M; (c) in TBAHFP
at the level of 0.1 M of LiCl, 6=100 mV s−1
.
3.4.3. Chloride medium
The pattern of the cyclic voltammograms of com-
plexes 5 and 6 are almost the same of that obtained
Acknowledgements
from the addition of LiCl described above.
.
We thank the Comunidad Auto´noma de Madrid
(Spain) for financial support (projects 06M/041/96 and
07M/0015/99). We also thank Professor J. Soria for the
use of the EPR spectrometer.
The reduction potential values of complexes 5 and 6
in LiCl are in the range of the Fe–SOD mimic observed
[17]. Therefore, these compounds might be SOD
mimetic.
3.4.4. Cobalt complexes
Voltammetric behaviour of cobalt complexes 7 and 8
was studied in the potential range from +1 to −2.5 V.
The electrochemical behaviour of both complexes is
very close according to the same geometry proposed
from the spectroscopic data for these complexes. Both
complexes exhibit waves only in the negative range of
potential. Cyclic voltammogram of 7 (Fig. 7) shows an
irreversible reduction wave at −1.85 V attributed to
the insaturations of the ligand. In addition, it exhibits a
cathodic wave at −0.97 with the anodic counterpart at
−0.91 V corresponding to a cobalt(II) centred reduc-
tion process [45]. Study of the dependence of the
voltammetric parameters with scan rate in the interval
25–800 mV s−1 show a lineal relationship of ipc with
61/2. The peak potential shows a small dependence with
the scan rate. The ratio ipa–ipc is close to unity. From
these data, it can be deduced that this redox couple is
References
[1] D.L. Klayman, J.P. Scovill, J.F. Bartosevich, J. Bruce, J. Med.
Chem. 26 (1983) 35 and references therein.
[2] J.P. Scovill, D.L. Klayman, C. Lambos, G.E. Childs, J.D.
Nortsch, J. Med. Chem. 27 (1984) 87.
[3] M.J.M. Campell, Coord. Chem. Rev. 15 (1975) 279.
[4] S. Padhye, G.B. Kauffman, Coord. Chem. Rev. 63 (1985) 127.
[5] D.X. West, S.B. Padhye, B.P. Sonawane, Struct. Bond. 76 (1991)
1.
[6] E.J. Blanz, F.A. French, Cancer Res. 28 (1968) 2419.
[7] S. Padhye, R. Chikate, A. Kumbhar, J.M. Shallon, M.P. Chitnis,
Biometals 5 (1992) 67.
[8] D.X. West, A.E. Liberta, S.B. Padhye, R.C. Chikate, B.P.
Sonawane, A.S. Kumbhar, R.G. Yerande, Coord. Chem. Rev.
123 (1993) 49.
[9] I. Haiduc, A. Silvestru, Coord. Chem. Rev. 99 (1990) 253.
[10] C.M. Liu, G. Xiong, X. You, Y. Liu, Polyhedron 15 (1996)
4565.
[11] Z. Durackova´, J. Labuda, J. Inorg. Biochem. 58 (1995) 297.

2068
M. Can˜adas et al. / Polyhedron 19 (2000) 2059–2068
[12] D. Zhang, D.H. Busch, P.L. Lennon, R.H. Weis, W.L. Neuman,
D.P. Riley, Inorg. Chem. 37 (1998) 956.
[29] E. Franco, E. Lo´pez-Torres, M.A. Mendiola, M.T. Sevilla,
Polyhedron 19 (2000) 441.
[30] A. Arquero, M. Can˜adas, M. Mart´ınez-Ripoll, M.A. Mendiola,
A. Rodr´ıguez, Tetrahedron 54 (1998) 11 271.
[31] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81.
[32] A.B.P. Lever, Inorganic Electronic Spectroscopy, second ed.,
Elsevier, Amsterdam, 1984.
[33] E. Colacio, E. Lo´pez-Torres, M.A. Mendiola, J.M. Moreno,
Proc. 9a Reunion Cientifica de Qu´ımica Inorgania, Bilbao
(Spain) 2000.
[13] M. Rodriguez, I. Morgenstern-Baradau, M. Cesario, J. Guilhem,
B. Keita, L. Nadjo, Inorg. Chem. 35 (1996) 7804.
[14] B. Kwak, K.W. Cho, M. Pyo, M.S. Lah, Inorg. Chim. Acta. 290
(1999) 21.
[15] M.S. Lah, M.M. Dixon, K.A. Pattridge, W.C. Stallings, J.A.
Fee, M.L. Ludwing, Biochemistry 34 (1995) 1546.
[16] Y. Nishida, I. Watanabe, K. Unoura, Chem. Lett. (1991)
1517.
[34] R. Aasa, J. Chem. Phys. 52 (1970) 3919.
[35] J. Costa, R. Delgado, M.G.D. Drew, V. Fe´lix, R.T. Henriques,
J.C. Waerenborgh, J. Chem. Soc., Dalton Trans. (1999) 3253.
[36] J. Garc´ıa-Tojal, J.L. Pizarro, L. Lezama, M.I. Arriortua, T.
Rojo, Inorg. Chim. Acta 278 (1998) 150.
[37] B. Castro, C. Freire, Inorg. Chem. 29 (1990) 5113.
[38] A.J. Blake, R.O. Gould, M.A. Halcrow, M. Schro¨eder, J. Chem.
Soc., Dalton Trans. (1993) 2909.
[39] E. Pereira, L. Go´mez, B. Castro, Inorg. Chim. Acta 271 (1998)
83.
[40] M.A.A.F.C.T. Carrondo, B. Castro, A.M. Coelho, D.
Domingues, C. Freire, J. Morais, Inorg. Chim. Acta 205 (1993)
157.
[17] C.K. Vance, A.-F. Miller, J. Am. Chem. Soc. 120 (1998) 461.
[18] J.R. Lancaster (Ed), The Bioinorganic Chemistry of Nickel,
VCH, New York, 1988.
[19] A.F. Kolodziej, Prog. Inorg. Chem. 41 (1994) 493.
[20] M. Beley, J.P. Collin, R. Rupert, J.P. Sauvage, J. Am. Chem.
Soc. 108 (1986) 7461.
[21] K. Bujno, R. Bilewicz, L. Siegfried, T. Kaden, J. Electroanal.
Chem. 407 (1996) 131.
[22] H. Aga, A. Amarata, Y. Hisaeda, J. Electroanal. Chem. 437
(1997) 111.
[23] E. Kimura, S. Wada, M. Shionoya, Y. Okazaki, Inorg. Chem. 33
(1994) 770.
[24] E. Mejeritskaia, F. Luo, C.A. Kelly, B. Koch, E.M. Gundlach,
E.L. Blinn, Inorg. Chim. Acta 246 (1996) 295.
[25] A. Berkessel, M. Bolte, T. Newmann, L. Seidel, Chem. Ber. 129
(1996) 1183.
[26] P. Souza, M.A. Mendiola, A. Arquero, V. Ferna´ndez, E. Gutie´r-
rez-Puebla, C. Ru´ız-Valero, Z. Naturforsch., Teil B 49 (1994)
263.
[41] F. Azevedo, M.A.A.F.C.T. Carrondo, B. Castro, M. Convery,
D. Domingues, C. Freire, M.T. Duarte, K. Nielsen, I.C. Santos,
Inorg. Chim. Acta 219 (1994) 47 and references therein.
[42] F.V. Lovecchio, E.S. Gore, D.H. Busch, J. Am. Chem. Soc. 96
(1974) 3109.
[43] A.M. Tait, F.V. Lovecchio, D.H. Busch, Inorg. Chem. 16 (1977)
2206.
[27] A. Arquero, M.A. Mendiola, P. Souza, M.T. Sevilla, Polyhedron
15 (1996) 1657.
[28] Z. Durackova´, M.A. Mendiola, M.T. Sevilla, A. Valent, Bioelec-
trochem. Bioenerg. 48 (1999) 109.
[44] V.M. Leovac, L.S. Jovanovic, V.I. Cesljevic, L. Bjelica, C.B.
Arion, N.V. Geberleu, Polyhedron 13 (1994) 3005.
[45] A.A. Isse, A. Gennaro, E. Vianello, J. Chem. Soc., Dalton
Trans. (1993) 2091.
.