Nickel(II) Complexes with Sulphonylhydrazone Derivatives: Spectroscopic and Electrochemical Studies
The synthesis and characterization of benzene and p-toluenesulfonylhydrazones
derived from salicylaldehyde and 2-hydroxyl-1-naphthaldehyde and their
Ni(II) complexes are reported. The structural investigation of these compounds
was based on elemental analysis, magnetic moment and spectral (ultraviolet,
infrared and 1H-NMR). The Ni(II) complexes were diamagnetic.
The stoichiometry of all the complexes was found to be 1:2 and the geometry
around the nickel ions is square planar. The electrochemical behavior
of the Ni(II) complexes was investigated in DMSO by cyclic voltammetry
(CV), rotating disc electrode (RDE) and coulometry. The complexes displayed
Ni(III)/Ni(II) couples irreversible waves and the substitution of the
phenyl by naphthyl fragments causes a negative shift in the formal potential.
It is known that the transition metal complexes play a central role in
the conduction of molecular materials, which display unusual conducting
and magnetic properties and find applicability in material chemistry,
supramolecular and biochemistry (Chandra and Kumar, 2004, 2005; Chandra
et al., 2006). The hydrazone derivatives as ligands for transition
metal ions constitute an important class of ligands which have been extensively
studied in coordination chemistry mainly due to their facile synthesis,
their photochromic effect (Wong and Bruscato, 1968; Zady et al.,
1975; Jacques, 1984; Becker and Chagneau, 1992; Bulanov et al.,
2002; Hadjoudis and Mavridis, 2004), their physico-chemical properties
and application in many important chemical process that include sensors,
non-linear optics, medicine and others (Armstrong et al., 2003;
Schmitt et al., 2003; Bakir et al., 2004). It is well established
that transition metals are readily susceptible to oxidation and reduction
because of their ability to easily change oxidation states. Thus, the
high oxidation state transition metal complexes are well known to be biologically
important and interesting because of their redox enzymes properties (Margerum
et al., 1975; Bour et al., 1971; Ruiz et al., 1997;
Cervera et al., 1998).
The redox properties include oxidation of the central metal ion with
ligands have been previously studied and reported by Larabi et al.
(2003). The redox potential of the Cu(II)/Cu(III) and Ni(II)/Ni(III) complexes
have been shown to be markedly affected by the nature of the chelating
ligand with the complexes (Larabi et al., 2003; Schmidt and Chmielewski,
In earlier research (Larabi et al., 2003) salicylaldehyde benzenesulphonyl-hydrazone
(SBSH), naphthaldehyde benzenesulphonylhydrazone (NBSH), salicylaldehyde
p-toluene-sulphonylhydrazone (STSH) and naphthaldehyde p-toluenesulphonyl-hydrazone
(NSTH) (Scheme 1) were synthesized and the structures
of their Cu(II) complexes were reported. The present study is concerned
with the elucidation of these ligands structures by the use of spectroscopic
methods (UV, IR and 1H-NMR). In addition, information about
the stereochemistry of their Ni(II) complexes has been obtained from spectral
measurements. It is worthwhile to note that little attention on the electrochemical
behavior of metal complexes derived from hydrazone derivatives was given.
In order to contribute in this area, it is studied extensively the electrochemical
properties of the Ni(III)/Ni(II) couple, which was affected by the donor
environments. So, the present study is also to investigate the capacity
of the chelating ligands under investigation to stabilize nickel(III)
complexes and to elucidate the mechanism of the oxidation process by means
cyclic voltammetry, rotating disc electrode and bulk electrolysis techniques.
|| Presentation of the ligands
MATERIALS AND METHODS
All the chemicals used for the preparation of the ligands were of BDH
quality. Conductivity measurements were carried out in DMF (ca. 10-3
mol L-1) using a Tacussel conductivity bridge model 75. Magnetic
susceptibilities were determined using a Johnson Matthey balance at room
temperature (25 °C) with Hg[Co(SCN)4] as standard. Perkin-Elmer
PE 938 and Pye Unicam model SP. 3-300 spectrophotometers were used to
record the IR spectra using KBr pellets and as Nujol mulls between CsI
plates. UV-visible spectra were recorded in Nujol, in acetone and chloroform
on a Perkin-Elmer model 550-S spectrometer. 1H-NMR spectra
were recorded using Bruker Ac 200 spectrophotometer at Strasbourg University
(France). Elemental analysis were carried out in the Micro-analytical
unit at Cairo University (Egypt). The electrochemical experiments were
carried out using a Trace-lab50 from Radiometer which includes a polarographic
analyzer (Pol 150), a polarographic stand (MDE 150) and trace Master 5
software. Cyclic voltammetry was performed using a conventional three
electrodes system. The working electrode was a pre-polished glassy carbon
(GC) disc of 3 mm diameter (Radiometer). Potentials are expressed versus
the Ag/AgCl (KCl 3 mol L-1) electrode separated from the test
solution by a salt bridge containing the solvent/supporting electrolyte.
The auxiliary electrode was a platinum wire. The RDE study was performed
using Radiometer model BM-EDI101 rotating disc electrode. The rotating
speed ω was regulated by an Asservitex model CTV101 from the Radiometer.
The following solutions were studied: 0.5, 1.0, 1.5, 2.0, 3.0, 3.5, 4.0
and 4.5 mmol L-1 of complexes in DMSO and 0.1 mol L-1
N(Et)4ClO4 as supporting electrolyte.
Caution: N(Et)4ClO4 is sensitive to shock
or heat. The RDE voltammograms were recorded in each solution, using a
scan rate of 5 mV sec-1 and rotating speeds ω of 25, 50,
75, 100, 150, 200, 300 and 400 rpm. In The CV measurements, scan rate
v of 10, 25, 50, 100, 200, 400, 500, 1000 and 2000 mV sec-1
were employed. All experiments were carried out at 25 °C ±
0.1 using a Julabo thermostat. In the coulometric experiments, the auxiliary
electrode was separated from the solution by a glass frit disk and the
working electrode was controlled by a Radiometer PGP201 potentiostat.
Synthesis of the ligands: Benzene and p-toluenesulphonylhydrazine
(BSH, TSH) were prepared according to literature procedures (Vogel, 1989)
and the ligands SBSH, NBSH, STSH and NTSH were synthesized as reported
earlier (Larabi et al., 2003).
Synthesis of the complexes: The metal complexes were prepared
using a general method. A hot absolute EtOH solution of Ni(II) acetate
(1 mmol) was added to a hot solution of corresponding ligand (2 mmol)
in EtOH with continuous stirring. The precipitated was filtered off hot,
washed several times with an absolute EtOH and dried in a dessicator over
RESULTS AND DISCUSSION
Structural studies: Analytical results and physical properties
of the ligands and the product complexes are given in Table
1. They are air-stable, insoluble in most common organic solvents
but easily soluble in DMF and DMSO. The molar conductivities of the complexes
in DMF (25 °C) are in the 3-7 ohm-1 cm2 mol-1
range, indicating a non-electrolytic nature (Geary, 1971).
The position of the significant IR bands of all ligands (H2L1,
H2L2, H2L3 and H2L4)
and their nickel(II) complexes are summarized in Table 2.
The IR spectra of H2L1 show two bands at 3020 and
2880 cm-1 assignable to γaOH and γsOH
vibrations, respectively. These bands are observed at 3420 and 2920 cm-1
for H2L3. The existence of those bands at lower
wave-numbers suggests the presence of intramolecular hydrogen bonding
of the type (O-H....N) (Bullock and Tajmir-Riahi, 1978). Also, the spectra
show strong band at 3160 for H2L1 and 3200 cm-1
for H2L3 assigned to γNH vibration.
|| Colors, melting point, partial elemental analysis and
molar conductivities of the metal complexes
||Infrared spectra of the ligands and their metal complexes
|HB: Hydrogen bonding, -: Absence of the bond
the bands at ~ 1620, ~ 1430 and ~ 1270 cm-1 are assigned to
γ (C=N), γ (C-O) and γOH vibrations, respectively. The
four bands at ~1325, ~ 1170, ~ 570 and ~ 480 cm-1 are attributed
to γsSO2, γasSO2,
δSO2, γSO2 vibrations, respectively and
remain more or less at the same positions as reported in literature (Bellamy,
1958). The observation of broad but weak bands in the 2000-1800 and 2400-2200
cm-1 regions suggests the existence of hydrogen bonding of
the type N-H....N (intermolecular hydrogen bonding). The IR spectra of
both ligands in CHCl3 show the obscure of hydrogen bonding.
The 1H-NMR spectra of H2L1 and H2L2 in d6-DMSO show two singlet signals at 11.50 and 11.10 ppm,
downfield of TMS, with equal ratio and assigned to the protons of OH and
NH groups, respectively. Those two signals disappear upon deuteration.
The IR spectra of H2L2 and H2L4 show three important bands. H2L2 exhibits bands
at 3180, 3030 and 2880 cm-1 while H2L4 show bands at 3200, 3460 and 3040 cm-1. These bands can be
assigned to γaOH γsOH and γNH vibrations,
respectively. The position of the latter bands suggests that the OH group
is strongly affected by strong intramolecular hydrogen bonding of the
type O-H....N (Bullock and Tajmir-Riahi, 1978). In addition, the observation
of broad weak bands in the 2000-1800 and 2400-2200 cm-1 regions
is also taken as evidence for the presence of hydrogen bonding. The 1H-NMR
spectra of H2L3 and H2L4 in
d6-DMSO show three singlet signals at 2.4, 10.2 and 11.5 ppm,
downfield of TMS, with ratio 3:1:1 and are assigned to the protons of
CH3, NH and OH groups, respectively. The latter two signals
disappear upon deuteration. All these observations suggest the structures
for the ligands given in Scheme 2.
The UV spectra of H2L1 and H2L3
in acetone show four bands. H2L1 exhibits the bands
at 438, 370, 330 and 304 nm while H2L3 shows the
bands at 430, 412, 380 and 340 nm. These bands are assigned to n →
π* (SO2), n → π* (C=N), π → π*
(SO2) and π → π* (C=N), respectively (Rao, 1975).
On the other hand, the UV spectra of H2L2 and H2L4
in acetone show five bands. H2L2 shows bands at
450, 418, 376, 360 and 330 nm while H2L4 exhibits
bands at 450, 438, 380, 360 and 328 nm. These bands may be assigned to
n → π* (SO2), n → π* (C=N), π →
π* (SO2) and π → π* (C=N) and π →
π* (naphtyl), respectively (Rao, 1975). It is interesting to point
that the above mentioned bands are split into two for each band in CHCl3.
This suggests that the ligands exist in two forms (free and hydrogen-bonded)
upon dissolving in CHCl3 as shown in Scheme 3.
In comparing the IR spectra of the Ni(II) complexes with the parent ligands
(Table 2), we observed that the ligands behave more
or less in the same way. The ligands coordinate in a bidentate manner
via the azomethin group (C=N) and the OH (phenolic or naphtolic) groups
forming six membered ring including the metal ions.
The displacement of a hydrogen atom from the OH group is proved as follows:
The pH drop and the conductance increase upon successive addition of
ligand solution on titrating against Ni(II) acetate solution prove the
liberation of ethanoic acid during complex formation. The spectral data
(IR and 1H-NMR) confirm participation of OH in bonding. Thus,
the γOH disappears, the C=N band shifts to lower wave number and
new bands appear in the 560-520 and 430-410 cm-1 regions assignable
to γ(M-O) and γ(M-N) (Ferraro, 1971; Chandra et al.,
2005), respectively. The latter result supports the involvement of nitrogen
in coordination. Finally the test of OH group by spot test technique (Feigl,
1982) is negative.
The diamagnetic behaviors as well as the observation of a broad band
centered at ca. 476 nm assigned to 1A1g → 1A2g
transition are evidences for square-planar geometry (Lever, 1968; Sacconi,
1966) around the nickel(II) ion. The disappearance of the OH proton and
the existence of the NH proton in the 1H-NMR spectra of the
nickel complexes, prove the replacement of a hydrogen atom from the OH
Electrochemical studies: The electrochemical behavior of all the
Ni(II) complexes are similar in the same conditions and depends on the
|| Presentation of the inter- and intramolecular hydrogen
|| Presentation of the two forms of the ligands
Cyclic voltammogram of 1 mmol L-1 [Ni(HL1)2]
in DMSO-TEAP (0.1 mol L-1) at GCE; scan rate 0.1 V sec-1
at 25 °C. Inset shows the CV of H2L1 (1
mmol L-1) in the same solution
The cyclic voltammogram of 1 mmol L-1 [Ni(HL1)2]
which is similar to that of [Ni(HL3)2], in DMSO
with 0.1 mol L-1 N(Et)4ClO4 as the supporting
electrolyte is shown in Fig. 1. The voltammogram obtained,
in the anodic direction, at a glassy carbon electrode shows two prominent
oxidation waves at Epa values 0.58 V (peak 1) and 1.22 V (peak
2). It should be mentioned that the supporting electrolyte N(Et)4ClO4 in DMSO did not show any redox activity in the potential range studied.
So The anodic process (peak 2) is a result of the redox process of the
corresponding ligand (H2L1). This signal is also
observed in the CV of H2L1 under similar conditions
and may be attributed to the irreversible oxidation of the NH group. The
other peak (peak 1) observed in the CV plot is assumed to be a result
of the oxidation occurring at the Ni(II) ion. This redox couple studied
in the interval of 10-2000 mV sec-1 shows a linear variation
of Ipa1 versus v1/2 at complex concentrations ≤
3 mmol L-1. Moreover the slope ΔE/Δlog v has the
value of 0.046 V that is larger those expected for reversible process.
These suggests a diffusion controlled irreversible Ni(II)-Ni(III) process
at complex concentrations ≤ 3 mmol L-1. The value of the
symmetry coefficient α was also determined using Eq.
1 (Nicholson and Shain, 1964) and was found to be 0.45 when n = 1,
which confirmed the irreversible nature of the electrode process Ni(II)-Ni(III).
where, Ep/2 is the half-peak potential, n is the total number
of electrons involved in the reaction.
On the other hand, the Fig. 2 shows that at higher
complex concentration and for v>400 mV sec-1 the increase
of Ipa1 with v1/2 was less than linear. These results
indicate that the complex oxidation occurs via an irreversible electron
transfer followed by a chemical reaction at complex concentrations >3
We also carried out the RDE experiments for [Ni(HL1)2]
at different concentrations in DMSO solutions in order to elucidate the
reaction mechanism. Figure 3 shows the results plotted
according to Levich equation:
where, D, v, ω and C0 are the diffusion coefficient,
the kinematic viscosity, the rotation speed and the bulk concentration
of the reactant in the solution respectively and all other parameters
have their conventional meanings.
Levich equation predicts that the plot of Il vs. ω1/2
should be linear.
Variation of the anodic peak currents, Ipa1,
versus v1/2 for various [Ni(HL1)2]
concentrations: ( ■ ) 0.5; (•) 1.0; ( ▲ ) 1.5; (
▼ ) 2.0; ( ♦ ) 3.0; ( ◄ ) 4.0 and ( ► ) 4.5
Levich plots of the first anodic limiting current of
[Ni(HL1)2] at different concentrations: ( ■
) 0.5; (•) 1.0; ( ▲ ) 1.5; ( ▼ ) 2; ( ♦ )
3; ( ◄ ) 4.0 and ( ► ) 4.5 mmol L-1 at a GCE;
V=10 mV sec-1
This was true only when the complex concentration was
≤ 3 mmol L-1 (Fig. 3), while for concentrations
higher than 3 mmol L-1 the corresponding plots were found to
be of curved shape. This result is an indication of a kinetic limitation
(Razmi-Nerbin and Pournaghi-Azar, 2002).
The behavior of the oxidation peak currents (Ipa1) of [Ni(HL3)2]
and the Il vs. ω1/2 equations is identical
to that reported for CV and RDE studies of [Ni(HL1)2].
The cyclic voltammogram of [Ni(HL2)2] at a glassy
carbon electrode with DMSO as the solvent (Fig. 4) is
very similar to that of [Ni(HL4)2]. The voltammogram
obtained, in the anodic direction, displays three oxidative peaks (1,
3 and 2) at 25 °C and at a sweep rate of 100 mV sec-1.
Cyclic voltammogram of 1 mmol L-1 [Ni(HL2)2]
in DMSO-TEAP (0.1 mol L-1) at GCE; scan rate 0.1 V sec-1
at 25 °C. Inset shows the CV of H2L2 (1
mmol L-1) in the same solution
||Cyclic voltammograms of 2 mmol L-1 [Ni(HL2)2]
at different sweep rates
The two anodic processes (peak 3 and 2) together with the cathodic peaks
(3`c and 3c) are a result of the redox processes of H2L2 and H2L4 ligands. These signals are also observed
in the CV of H2L2 and H2L4 under similar conditions. Not that the reverse scan at 1100 mV rises the
appearance of peaks 3`c and c indicating the association of these
peaks and peak 3. The peaks 3 and 2 are tentatively attributed to the
irreversible oxidation of the naphtyl group and NH group, respectively.
The peak 1 may be assigned to the Ni(III)/Ni(II) redox couple.
On the other hand, we can note that on increasing the scan rate, for
[Ni(HL2)2] and [Ni(HL4)2]
at 2 mmol L-1, from 100 to 2000 mV sec-1 (Fig.
5), the oxidation-reduction waves become more reversible-like. Probably
this change is due to increase of scan rate and diffusion problems occur
(Perez et al., 2005). This may be also due to the fact that the
electron transfer reaction is followed by a chemical reaction. Note that
the [Ni(HL1)2] and [Ni(HL3)2]
complexes do not show the same behavior over the range of studied voltage
The peak current Ip, for oxidation peak 1 increased linearly with
bulk solution concentration of [Ni(HL2)2] over the concentration
range 0.5-2 mmol L-1. At concentrations higher than 2 mmol L-1
this increase was less than linear. Furthermore, it was shown that Ip
increases linearly as a function of square root of the voltage sweep rate over
the sweep rate range 100-2000 mV sec-1 at complex concentrations
≤ 2 mmol L-1. At higher concentrations, the current increases with
increasing square root of the voltage sweep rate, but was found to be non-linear.
Moreover, the RDE study shows that for complex concentrations ≤ 2 mmol L-1,
the limiting current increases linearly with increasing electrode rotation speed.
At higher complex concentrations, this increase was found to be of curved shape.
These behaviors indicate that the peak 1 oxidation of [Ni(HL2)2]
is a diffusion controlled reaction over the entire range of voltage sweep rate
studied at complex concentrations ≤ 2 mmol L-1. At higher complex
concentrations (>2 mmol L-1), the current is governed by the rate
of the charge transfer or of a chemical reaction. Moreover, it is possible to
suggest that the formation of Ni(III) complexes undergoes an electron transfer
followed by relatively fast chemical reaction, which probably results in the
partial dissociation of the complexes.
On the other hand, a controlled-potential preparative electrolysis was
carried out in a divided cell. An exhaustive oxidation of [Ni(HL2)2]
was performed at constant potential (500 mV). The anolyte was 20 mL of
0.5 mmol L-1 [Ni(HL2)2] solution. A Pt
electrode was used as the cathode. The result indicates that approximately
1 electron was transferred per molecule (n = 1.2). Since the peak currents
Ipa1 for all the complexes studied have approximately the same
value, the number of electron involved in the oxidation of Ni(II) is 1
(for all the complexes). After exhaustive oxidation of [Ni(HL2)2],
the resulting solution was scanned from 0 to 1400 mV, the oxidized complex
did not exhibit any oxidation signal Ni(II) → Ni(III) as shown in
The electrochemical data of 2 mM nickel complexes studied are summarized
in Table 3. A change in the formal potentials of the
nickel(III)/nickel(II) couples, which occurs upon ligand substitution
along this series is observed. A comparison of the redox potentials of
the [Ni(HL3)2] and [Ni(HL4)2]
complexes where both ligands H2L3 and H2L4
have the same set of donor atoms but differ in the size of central chelating
agents shows that the nickel(III) complex [Ni(HL4)2]
is more stable thermodynamically, i.e., the Ni(III)/Ni(II) potential is
lowered by 140 mV. The same result was obtained by comparing the redox
potentials of [Ni(HL1)2] and [Ni(HL2)2].
The Ni(III)/Ni(II) potential is, here, lowered by 150 mV.
||Cyclic voltammograms of 0.5 mmol L-1 [Ni(HL2)2]
before (A) and after (B) exhaustive oxidation
|| Cyclic voltammetric data of the studied complexes*
*c = 2 mmol L-1 in DMSO solution (0.1 mol
L-1 N(Et)4ClO4), E values (V versus
Ag/AgCl/3 mol L-1 KCl), Scan rate 100 mV sec-1,
Epa and Epc are the anodic and cathodic peak
In this study, we could synthesize Ni(II) hydrazone derivatives complexes
containing nitrogen and oxygen donor atoms. The structure determinations
of ligands and their complexes were established by elemental analysis,
magnetic moments and UV, IR and 1H-NMR spectra. In all complexes,
the geometry around the nickel(II) ions is square planar. It was suggested
that the ligands exist in two forms (free and hydrogen-bonded). Electrochemistry
data show that the Ni(II) complexes displayed Ni(III)/Ni(II) couples irreversible
waves and that the substitution of the phenyl by naphtyl fragments causes
a large enough increase in aromatic character to produce a negative shift
in the formal potential.
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