Spectroscopic and biological activities studies of bivalent transition metal complexes of Schiff bases derived from condensation of 1,4-phenylenediamine and benzopyrone derivatives

Article abstract of DOI:10.1016/j.saa.2013.08.037

Many tools of analysis such as elemental analyses, infrared, ultraviolet-visible, electron spin resonance (ESR) and thermal analysis, as well as conductivity and magnetic susceptibility measurements were used to elucidate the structures of the newly prepared Co(II), Ni(II) and Cu(II) complexes with Schiff bases derived from the condensation of 1,4-phenylenediamine with 6-formyl-7-hydroxy-5-methoxy-2-methylbenzo-pyran-4-one (H₂L) or 5,7-dihydroxy-6-formyl-2-methylbenzopyran-4-one (H₄L). The data showed that all formed complexes are 1:1 or 2:2 (M:L) and non-electrolyte chelates. The Co(II) and Cu(II) complexes of the two Schiff bases were screened for antibacterial activities by the disk diffusion method. The antibacterial activity was screened using Escherichia coli and Staphylococcus capitis but the antifungal activity was examined by using Aspergillus flavus and Candida albicans. The Results showed that the tested complexes have antibacterial, except CuH₄L, but not antifungal activities.

Full text of DOI:10.1016/j.saa.2013.08.037

Accepted Manuscript
Spectroscopic and biological activities studies of bivalent transition metal com‐
plexes of Schiff bases derived from condensation of 1,4-phenylenediamine and
benzopyrone derivatives
Omaima E. Sherif, Nora S. Abdel-Kader
PII:
S1386-1425(13)00911-6
http://dx.doi.org/10.1016/j.saa.2013.08.037
SAA 10896
DOI:
Reference:
To appear in:
Spectrochimica Acta Part A: Molecular and Biomo‐
lecular Spectroscopy
Received Date:
Revised Date:
Accepted Date:
7 April 2013
25 July 2013
2 August 2013
Please cite this article as: O.E. Sherif, N.S. Abdel-Kader, Spectroscopic and biological activities studies of bivalent
transition metal complexes of Schiff bases derived from condensation of 1,4-phenylenediamine and benzopyrone
derivatives, Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy (2013), doi: http://dx.doi.org/
10.1016/j.saa.2013.08.037
This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers
we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and
review of the resulting proof before it is published in its final form. Please note that during the production process
errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

Spectroscopic and biological activities studies of bivalent transition
metal complexes of Schiff bases derived from condensation of 1,4-
phenylenediamine and benzopyrone derivatives
Omaima E. Sherif and Nora S. Abdel-Kader
Chemistry Department, Faculty of Science, Cairo University, Giza, Egypt
Abstract
Many tools of analysis such as elemental analyses, infrared, ultraviolet-visible,
electron spin resonance (ESR) and thermal analysis, as well as conductivity and
magnetic susceptibility measurements were used to elucidate the structures of the
newly prepared Co(II), Ni(II) and Cu(II) complexes with Schiff bases derived from
the condensation of 1,4-phenylenediamine with 6-formyl-7-hydroxy-5-methoxy-2-
methylbenzo-pyran-4-one (H2L) or 5,7-dihydroxy-6-formyl-2-methylbenzopyran-4-
one (H4L). The data showed that all formed complexes are 1:1 or 2:2 (M:L) and non-
electrolyte chelates. The Co(II) and Cu(II) complexes of the two Schiff bases were
screened for antibacterial activities by the disc diffusion method. The antibacterial
activity was screened using Escherichia coli and Staphylococcus capitis but the
antifungal activity was examined by using Aspergillus flavus and Candida albicans.
The Results showed that the tested complexes have antibacterial, except Cu-H4L, but
not antifungal activities.
Keywords: Schiff bases, transition metal complexes, biological activities
1. Introduction
Salen-type ligands, one of the oldest classes of ligands in coordination
chemistry, have been used extensively to complex transition metals [1]. The majority
of salen ligands reported in the literature are symmetric and available as their N,N'-
disubstituted derivatives [2]. Salen-type ligands have been studied widely especially
in regard to their use in catalytic synthesis [3, 4]. In addition to, the C=N linkage in
azomethine derivatives is an essential structural requirement for biological activity.
Several azomethines have been reported to possess remarkable antibacterial [5-9],
antifungal [10-12], anticancer [13-16] and diuretic activities [17]. In view of this, we
have designed a derivatives of salen type Schiff bases and their complexes. Quite
recently, the Schiff bases derived from the condensation of 6-formyl-7-hydroxy-5-
1

methoxy-2-methylbenzopyran-4-one and 5,7-dihydroxy-6-formyl-2-methylbenzopyran-4-
one with aliphatic diamines (ethylenediamine and trimethylenediamine) and their
complexes were previously prepared and characterized [18-21]. In this paper, we
report the synthesis and characterization of new transition metal complexes of Co(II),
Ni(II) and Cu(II) with Schiff bases derived from condensation of 6-formyl-7-hydroxy-
5-methoxy-2-methylbenzopyran-4- one or 5,7-dihydroxy-6-formyl-2-methylbenzopyran-
4-one and 1,4-phenylenediamine. The complexes have been characterized by several
tools of analyses such as elemental analyses, magnetic moment measurements, molar
conductance, ESR, IR, UV-Vis spectra as well as thermal gravimetric analysis (TG)
have also been studied for some selected metal complexes. The cobalt and copper
complexes were assayed by the disc diffusion method for antibacterial and antifungal
activity and compared with that of free Schiff bases.
2. Experimental
2.1 Materials and reagents
All chemicals used in this investigation were chemically pure grade (BDH or
Aldrich). They included nickel (II) chloride, cobalt (II) acetate, cobalt (II) perchlorate,
copper (II) acetate, 1,4-phenylenediamine, 5-methoxy-2-methyl-furanobenzopyran-4-
one, potassium dichromate, sulphoric acid, hydrochloric acid. All organic solvents
used in this study were either purified by recommended methods [22] or obtained as
grade materials from BDH. Doubly distilled H2O was always used.
2.2. Analysis and physical measurements
Carbon, hydrogen and nitrogen were analyzed by standard microanalysis
methods at Microanalytical center, Cairo University, Giza, Egypt. UV-Vis spectra of
the metal complexes in DMF were recorded on UV/Vis-NIR 3101 PC Shimadzu
spectrophotometer. IR spectra of the ligands and their metal complexes, as KBr discs,
were recorded on a Shimadzu FTIR spectrometer. Molar conductivity of 10-3mol/L
solutions of the complexes in DMF was measured on the conductivity meter ORION
model 150 of 0.6 cell constant. The ESR spectra were recorded for Cu-H2L and Cu-
H4L on a Bruker EMX spectrometer working with frequency 9.715 GHz and magnetic
field set of 3371 E. Thermal analysis has been carried out using Shimadzu-50
o
Thermal Analyzer from room temperature to 600 C at a heating rate 10oC min-1.
Analysis of the metal ions were carried out by dissolving the complexes in
2

concentrated nitric acid followed by hydrogen peroxide, neutralizing the diluted
aqueous solutions with sodium hydroxide to pH 5.5 and the metal content was
determined by recommended method [23] or were determined by the weight of the
complex residue after thermal decomposition. Magnetic susceptibilities of the chelates
were measured at room temperature using a magnetic susceptibility Cambridge
England Sherwood Scientific. The effective magnetic moments were calculated from
the expression µeff = 2.828(XM T)1/2 B.M., where XM is the molar susceptibility
corrected using Parcel’s constants for diamagnetism of all atoms in the compounds [24].
2.3. Synthesis
2.3.1. Synthesis of the Schiff-bases
The Schiff base H2L was prepared by adding 5 mmole of 20 mL ethanolic
solution of 1,4-phenylenediamine gradually with continuous stirring to a solution of
10 mmole of 6-formyl-7-hydroxy-5-methoxy-2-methybenzopyran-4-one which
prepared from 5-methoxy-2-methylfuranobenzopyran-4-one as previously mentioned
[25]. While, the Schiff base H4L was prepared by addition of 1,4-phenylenediamine
(5 mmole) in 20 ml ethanol dropwise with continuous stirring to 10 mmole of
ethanolic solution of 5,7-dihydroxy-6-formyl-2-methylbenzopyran-4-one (prepared as
previously mentioned [25]). The reaction mixtures were stirred at room temperature
for at least 30 min and the solids obtained were filtered off, washed with ethanol and
recrystallized from dimethylformamide. Scheme 1 illustrates the synthesis of the
Schiff bases H2L and H4L.
3

OCH3
O
OCH3
O
OH
O
OHC
OHC
HO
HCl/H O
2
H SO / K Cr O
2
4
2
o
2
7
70-80
C
Reflux
O
O
CH3
O
O
HO
CH3
CH3
5-Methoxy-2-methyl-
furanobenzopyran-4-one
6-Formyl-7-hydroxy-5-methoxy-
2-methylbenzopyran-4-one
5,7-dihydroxy-6-formyl-
2-methybenzopyran-4-one
X
O
OHC
2
O
HO
CH3
p-phenylenediamine
H
C
X
O
6'
7'
4'
3'
5'
N
H
O
X
5
H
2'
O
O
CH3
C
6
1'
4
3
N
H
2
7
H3C
O
1
O
X = OCH for Schiff base H L
3
2
= OH for Schiff base H L
4
Scheme1. Preparation of the Schiff bases
2.3.2. Synthesis of the metal complexes
The solid chelates were prepared by mixing a hot ethanolic solution of
Cu(CH3COO)2.H2O, NiCl2.6H2O, Co(CH3COO)2.6H2O and Co(ClO4)2.6H2O with the
Schiff bases sufficient to form 1:1, 1:2 and 2:1 (M:L) complexes. The reaction
mixture was stirred on a water bath. The solid chelates were filtered off and washed
several times with second distilled water and ethanol until the filtrate becomes
colorless. The obtained complexes were kept in a vacuum desiccator and subjected to
elemental analysis.
3. Results and discussion
Elemental analyses and some physical properties of the isolated pure
complexes are listed in Table 1. The analytical results demonstrate that all the
4

prepared complexes have 1:1 or 2:2 (M:L) stoichiometry. The molar conductance
measurements in DMF of all the 1:1 (M:L) chelates are in the range 11.7-34.8 ohm-1
cm2 mole-1. These values are lower than those expected for an electrolyte. These
observations indicate a non-electrolytic nature of all complexes [26] in DMF at room
temperature. The solid complexes are soluble in DMF and DMSO. The structures and
properties were also characterized by FT-IR, UV-Vis, ESR spectroscopy, thermal
analysis (TGA) and magnetic moment measurements.
3.1. IR spectra
The most important bands occurring in the IR spectra of the free ligands and
their complexes have been compared in order to ascertain the bonding sites in the
chelating ligands, Table 2. The IR spectra of H2L and H4L ligands show a broad band
at 3433 and 3430 cm-1 due to phenolic OH stretching, respectively. The bands at 1658
and 1340 cm-1 are assigned as ν(C=O) of benzopyrone carbonyl in position four and
ν(C-N) vibration, respectively. The stretching vibration for the (C=N) (azomethine) in
the free Schiff bases H2L and H4L was observed at 1608 and 1632 cm-1, respectively.
H2L and H4L ligands display bands at 1167 and 1170 cm-1 which can be assigned to C-
OH stretching vibration. On the complexation, the band due to ν(C=N) in the IR
spectrum of H2L Schiff base exhibit an upward shift by about 9cm-1 in case of Coa-H2L.
On other hand, this band disappeared in the IR spectra of the complexes of Cu-H2L,
Cop-H2L and Ni-H2L complexes. The ν(C=N) stretching shows upward shift by about
12-15 cm-1 in the IR spectra of the complexes of H4L. This reveals that the azomethine
group is one of the coordination sites of the Schiff base ligands H2L and H4L [27]. This
is also confirmed by the appearance of bands in the range 473-486 cm-1 due to the
formation of the ν (M→N). The shifting towards higher value for the band due to C-OH
stretching vibration in the spectra of complexes confirms the participation of the oxygen
anion of the hydroxyl group in position seven in the C-O-M bond. This mode of
bonding is supported by the appearance of new absorption band at 547-559 cm-1, which
is not observed in the spectrum of the free ligands. This band is assignable to ν(M-O) [28].
The band due to the stretching vibration of the carbonyl group in position four of
benzopyrane-4-one moiety shows no significant change in all the complexes indicating
that the C=O group did not participate in the chelate formation. The IR spectra of all the
complexes show new absorption bands in the region 3375-3416 cm-1 indicating the
presence of the water or ethanol molecules inside the coordination sphere of the
5

complexes. In addition, the band at 837-830 cm-1 in the IR spectra of complexes due to
the OH bending of the coordinated water molecules suggests that water molecules are
coordinated to the metal ions [28]. Moreover, strong evidence for the presence or
absence of water of crystallization and / or coordinated water supported by the
thermogram of the complexes. The two bonding sites in case of H2L and H4L Schiff
bases are the oxygen of the deprotonated hydroxyl groups of benzoprone moiety in
position seven and the azomethine nitrogen leading to a stable six membered chelate
ring. Charge balance for the divalent positive transition metal ions can be achieved by
the participation of OH- in the coordination sphere of the metal ion from solution in all
complexes except Cop-H2L complex. In this complex chloride ion could be detected by
elemental analysis and the IR spectrum is indicative of the presence of coordinated
perchlorate ion. The presence of a coordinated perchlorate group in the IR spectrum of
Cop-H2L complex at group is shown by the appearance of bands at 1160, 1035, 945-and
648 cm-1 in the IR spectra [29]. The studies reveal that the bis-Schiff bases H2L and
H4L behave as monobasic bidentate chelating agents through the oxygen of the
deprotonated hydroxyl groups in position seven of the benzopyrone moieties and
azomethine nitrogen atoms.
3.2. Electronic spectra and magnetic moment studies
The electronic absorption spectra of Schiff bases and their complexes are shown in
Figs. 1 and 2. Due to the limited solubility of the complexes in most organic solvents,
the electronic spectra are scanned using a saturated solution of dimethylformamide
(DMF). The electronic absorption spectra of the two ligands (H2L and H4L) exhibit
shoulder at 30769 cm-1 due to the - * transition within the azomethine (CH=N) group.
This shoulder disappeared in the spectra of complexes as a result of the participation
of azomethine nitrogens in coordination. The band appears at 25316 and 25510 cm-1
in the electronic spectra of H2L and H4L, respectively, is assigned to an
intramolecular charge transfer (CT transition) involving the whole molecule. The
electronic spectra of all complexes show bands in the range of 25510-27027 cm-1,
Table 3, which may be due to a LMCT transition. The measured μeff values of Cu-H4L
is 1.74 BM, this value suggests the square-planar geometry [30]. On the other hand,
Cu-H2L complex shows magnetic value of 2.21 B. M., this value is higher than the
spin-only value of 1.73 B. M. for one unpaired electron and indicates that this
complex is dimeric. In addition, Ni(II) forms complexes with H2L and H4L ligands
6

with different geometries. The observed diamagnetism of Ni-H2L suggests its square-
planar geometry but the magnetic moment of Ni-H4L complex (3.2 M. B.) lies within
the region expected for octahedral geometry. The room temperature magnetic
susceptibility measurement shows that the μeff values are 2.49 and 2.89 BM for Cop-
H2L and Cop-H4L respectively. These values indicate the presence of one unpaired
electron for square planar stereochemistry around Co(II) d7 complexes [31]. Co(II)
acetate with H2L and H4L complexes show unusual magnetic moment values (6.36
and 6.07 BM) which could be considered as 2:2 (M:L) complexes where two ligand
molecules are coordinated to two metal ions.
3.3. ESR spectra
To obtain further information about the stereochemistry and the metal-ligand bonding
and to determine the magnetic interaction in the metal complexes, the powder ESR
spectra of Cu-H2L and Cu-H4L complexes were recorded at room temperature with
frequency 9.715 GHz and magnetic field set of 3371 G. The spin-Hamiltonian
parameters of the two complexes were calculated, Table 4. The solid state ESR
spectra of the two complexes exhibit axially symmetric g-tensar parameters with g|| >
gꢀ > 2.003 indicating that the copper site has a dx2-dy2 ground state characteristic of
tetradedral, square planar or octahedral stereochemistry [32]. The values of gav were
calculated according to the equation gav = 1/3 (g|| + 2gꢀ) and the results are given in
Table 4. The g|| value is an important function for indicating the covalent character of
M-L bond. Kivelson and Neiman have reported that g|| value < 2.3 for covalent
character of metal-ligand band and g|| value > 2.3 for ionic character. For the two
copper complexes, the g|| is less than 2.3 revealing an appreciable covalent character
for M-L bond [33]. Based on IR, electronic spectra and ESR data, Cu-H2L complex
has octahedral structure but Cu-H4L complex has square planar structure. In the axial
7

symmetry, the g-values are related by the expression G = (g|| - 2.0023) / (g -2.0023)
where G is the exchange interaction parameter. According to Hathway [34], if the
value of G is greater than 4, the exchange interaction is negligible in the solid state,
whereas, when is less than 4, considerable exchange interaction is indicated in the
solid state complex. A forbidden magnetic dipolar transition for Cu-H2L complex is
observed at half-field (c.1500 G, g = 4.0) but the intensity is very weak. Its spectrum
is similar to that of binuclear Cu(II) complexes [35]. The appearance of a half-field
signal confirms the magnetic interaction between the two Cu(II) ions in 2:2 (M:L) Cu-
H2L complex; this is not observed in Cu-H4L complex, which is mononuclear.
Superhyperfine structure for this complex is not observed at a high-field excluding
any interaction of the nuclear spins of the nitrogen (I = 1) with the unpaired electron
density on Cu(II). The exchange interaction parameter, G, in Cu-H2L complex is less
than four (3.3) , suggesting that a copper-copper exchange interaction exist. The low
magnetic moment, 2.21 B. M., for the binuclear Cu-H2L complex indicates the
presence of strong exchange interaction in the solid state. The G value for Cu-H4L
complex is greater than four (4.15) suggests the absence of exchange coupling
between Cu(II) center in the solid state [32].
3.4. Thermogravimetric analysis (TG) of the solid complexes
The results of the thermal analysis of the complexes are given in Table 5.
Thermal studies have been carried out using a thermogravimetric TG technique. The
experimental mass losses were calculated from the TG curves and the calculated ones
were obtained from the suggested formula based on the obtained data of the elemental
analyses. Fig. 3 illustrates the thermal behavior of the some chelates. All complexes
were thermally decomposed within the temperature 25-600 oC.
The TG curve of Coa-H2L complex shows the first degradation step within the
o
temperature range 50-130 C with an estimated mass loss of 5.56% (calc. 5.69%)
corresponding to the loss of four molecules of lattice water. But the TG curve of Coa-H4L
o
displays 8.33% (calc. 8.67%) mass loss in the temperature range 50-180 C which
corresponding to the loss of six water molecules of hydration. Thermal degradation of
the organic molecule of the two complexes starts above 240 oC. Further heating of the
two complexes result in continuous mass-loss till a constant mass where the CoO residue
is formed as a final product.
8

The TG curve of Cop-H2L chelate exhibits 5.00% mass loss in the temperature
o
range 50-130 C, which correlated to two water molecules outside the coordination
o
sphere. The mass loss in the range 130-330 C is attributed to the loss of one
coordinated ethanol molecule. Thermal degradation of the organic molecule starts
above 340 oC. Complete decomposition of the organic molecule ending at 595 oC with
formation of CoO as a final product.
For Ni-H4L chelate the mass loss occurring at 8.00% (calc. 8.07%)
corresponds to the loss of three coordinated water molecules in the temperature range
o
150-300 C. The organic part of the complex decomposes in two steps with the
formation of intermediates which finally decompose to the stable NiO at 597 oC.
3.5. Kinetics of thermal decomposition of the complexes
The kinetic and thermodynamic parameters of the thermal degradation process
were calculated from the integral method proposed by Coats-Redfern model [36]. The
Horowitz and Metzger [37] equation Cs = (n)1/1-n, where Cs is the weight fraction of
the substance present at DTG peak temperature, Ts, was used for the determination of
the value of the reaction order, and given by
Where: ms is the remaining weight at Ts, mo and
are the initial and final
weights of the substance, respectively.
Coats-Redfern equation for n = 1:
Where β is the heating rate, x is the fraction decomposed, A is Arrhenius pre-
exponential, R is the universal gas constant, T is the absolute temperature and Ea is
the activation energy. The term
is practically constant and the value of
vs. gave a straight line with a
≈ 1. Hence, a plot of
slope equal
equal
in which activation energy can be calculated and an intercept
in which Arrhenius pre-exponential factor can be also calculated.
The estimated values of Cs for the thermal decomposition of the complexes
were found in the range of 0.35-0.37 (Table 6). Thus indicates that the decomposition
of a series of H2L and H4L complexes was suggested to be first order, n=1, [38]. So
9

the values of the activation energy (Ea), Arrhenius constant (A), the activation entropy
(S*), the activation enthalpy (H*) and the free energy of activation (G*) are calculated
by applying Coats-Redfern equation for n =1 and given in the Table 6 where:
H* = E-RT
S* = Rln[Ah/KT]
G* = H*-TS*
From the obtained results, it is apparent that G values of the complex acquire
highly positive magnitudes. The high value of the energy of activation of the
complexes revealed the high stability of the investigated complexes due to their
covalent character [39]. The positive sign of ΔG revealed that the free energy of the
final residue is higher than that of the initial compound and all the decomposition
steps are non-spontaneous processes. The lowest value for ΔG of Cop-H2L indicates
that this ion has the highest tendency for interaction with the ligand. The enthalpy
values are positive for the investigated complexes supporting that a high temperature
is favourable for the complexation and the process is endothermic. The negative
values of ΔS for the degradation process indicates that more ordered activated
complex than the reactants and the decomposition reaction is slow [40]. The
correlation coefficient of the Arrhenius plots of the thermal decomposition steps were
found to be in the range 0.98-0.99 which indicates good fitness of the linear function.
4. Antimicrobial activity of some complexes
In azomethine derivatives, the C=N linkage is an essential structural
requirement for biological activity. The free Schiff bases H2L and H4Lwere screened
for antibacterial activities by micro-plate assay technique [25]. Escherichia coli and
Staphylococcus capitis were exposed to different concentrations of the Schiff bases.
As mentioned previously [25] the antibacterial effect of these Schiff bases on Gram-
negative were higher than that on Gram-positive bacteria moreover, the Schiff base
H2L which contain OCH3 on position five has higher antibacterial activity than H4L
that contain OH group on the same position [25]. These Schiff bases have no
antifungal activity. In the present study, the antimicrobial activities Cu(II) and
Co(II) acetate complexes of these two Schiff bases were carried out in order to
enhance their antimicrobial effect according to the conventional agar diffusion
method by tetracycline as a standard [41,42] using Staphylococcus aureus as Gram-
positive and Escherichia coli as Gram-negative bacteria, Aspergillus flavus (aflatoxin
10

producing fungus) and Candida albicans. The antibacterial inhibition zone values are
summarized in Table 7. As shown in the table the tested complexes have no
antifungal activity. Coa-H2L complex has the highest antibacterial activity, on the
other hand, Cu-H4L complex don’t have any bacterial activity on the tested
microorganisms. So the complexation lowered the efficiency of the Schiff bases as
antibacterial agents where the free Schiff bases have higher antibacterial effect [25] than
the complexes.
5. Conclusion
Cu(II), Ni(II) and two different salts of Co(II), cobalt (II) acetate and cobalt (II)
perchlorate, metal complexes of Schiff bases derived from the condensation of 1,4-
phenylenediamine with 6-formyl-7-hydroxy-5-methoxy-2-methylbenzopyran-4-one
(H2L) or 5,7-dihydroxy-6-formyl-2-methylbenzopyran-4-one (H4L) were prepared
and characterized by different spectroscopic techniques and TG and molar
conductance and magnetic moment measurements. The metal complexes exhibited
different geometrical arrangements such as square planar and octahedral
arrangements. In case Cu-H2L, Coa-H2L and Coa-H4L, two ligand molecules are
coordinated to two metal ions (2:2; M:L). The antimicrobial activities Cu(II) and
Co(II) acetate complexes of the two Schiff bases were carried out. The tested
complexes have no antifungal activity. Coa-H2L complex has the highest antibacterial
activity, on the other hand, Cu-H4L complex don’t have any bacterial activity on the
tested microorganisms.
The geometric structure, schemes 2-4, of the prepared complexes can be shown as
follows:
11

OCH3
OCH3
H
C
H
C
N
N
OH
OH
O
2
O
O
H O
2
OCH3
Cu
OCH3
Cu
O
2
H O
2
N
N
C
H
C
H
Scheme 2. The proposed structure of Cu-H2L complex
X
X
H
C
HC
N
N
O
. n H O
O
2
Co
O
N
N
O
C
CH
H
X
X
Complex
Coa-H2L
Coa-H4L
X
n
OCH3
OH
4
6
Scheme 3. The proposed structure of Coa-H2L and Coa-H4Lcomplexes
12

H
C
X
N
. n H O
X
H
C
2
HO
N
Y
O
OH
M
Y
OH2
Complex
Ni-H2L
Ni-H4L
Cu-H4L
Cop-H4L
X
OCH3
OH
Y
n
0
0
0
2
-
H2O
OH
-
-
OH
Scheme 4. The proposed structure of Ni-H2L, Ni-H4L, Cu-H4L and Cop-H4L and
complexes
13

6. References
[1] M.D. Hobday, T.D. Smith, Coord. Chem. Rev. 9 (1973), 311-337
[2] P.G. Cozzi, Chem. Soc. Rev. 33 (2004) 410-421.
[3] M. Calligaris, G. Nardin, C. Randaccio, Coord. Chem. 7 (1972) 385-389.
[4] H. Tye, J. Chem. Soc., Perkin Trans. 1 (2000) 275-298; Y.N. Ito, T. Katsuki, Bull.
Chem. Soc. Jpn. 72 (1999) 603-619; L. Canali, D.C. Sherrington, Chem. Soc.
Rev. 28 (1999) 85-93; M. Bandini, P.G. Cozzi, A. Umani-Ronchi, Chem.
Commun. (2002) 919-927; T. Katsuki, Synlett (2003) 281-297.
[5] S.V. More, D.V. Dongarkhadekar, R.N. Chavan, W.N. Jadhav, S.R. Bhusare, R.P.
Pawar, J. Indian Chem. Soc. 79 (2002) 768-769.
[6] A.K. Bhendkar, K. Vijay, A.W. Raut, Acta Ciencia Indica Chem. 30 (2004) 29-
32.
[7] Y.K. Vaghasiya, R.S. Nair, M. Baluja, S.S. Chanda, J. Serb. Chem. Soc., 69 (2004)
991-998.
[8] R. Fioravanti, M. Biava, G.C. Porretta, C. Landolfi, N. Simonetti, A. Villa, E. Conte,
A. Porta-Puglia, Eur. J. Med. Chem. 30 (1995) 123–132.
[9] J. Rhodes, H. Chen, S.R. Hall, J.E. Beesley, D.C. Jenkins, P. Collinsand, B.
Zheng, Nature 377 (1995) 71-75.
[10] H.M. Safwat, F.A. Ragab, N.M. Eid, G.M. Abdel, Egypt. J. Pharm. Sci. 29
(1988) 99-110.
[11] R. Mtrei, M. Yadawe, S.A. Patil, Orient. J. Chem. 12 (1996) 101-102.
[12] M.E. Hossain, M.N. Allam, J. Begum, M.A. Akbar, M.N. Uddin, F.E. Smith,
R.C. Hynes, Inorg. Chim. Acta 249 (1996) 207-213.
[13] K.P. Sharma, V.S. Jolly, P. Phatak,. Ultra Scient. Phys. Sci. 10 (1998) 263-
266.
[14] V.E. Kuz'min, A.G. Artemenko, R.N. Lozytska, A.S. Fedtchouk, V.P.
Lozitsky, E.N. Muratov, A.K. Mescheriakov, SAR QSAR Environ. Res. 16
(2005) 219-230.
[15] V. Barve, F. Ahmed, S. Adsule, S. Banerjee, S. Kulkarni, P. Katiyar, C.E.
Anson, A.K. Powell, S. Padhye, F.H. Sarkar, J. Med. Chem. 49 (2006) 3800-3808.
[16] D.R. Shkawat, S.S. Sabnis, C.V. Deliwala, Bull. Haffkine Inst. 1 (1973) 35-
39.
[17] C.T. Supuran, M. Barboiu, C. Luca, E. Pop, M. E.Brewster, A. Dinculescu,
Eur. J. Med. Chem. 31 (1996) 597-606.
14

[18] H.M. Abdel-Fattah, A.L. El-Ansary, N.S. Abdel-Kader, J Therm Anal
Calorim. 96 (2009) 961-969.
[19] A.L. El-Ansary, H.M. Abdel-Fattah, N.S. Abdel-Kader, J. Coord. Chem. 61
(2008) 2950-2960.
[20] A.L. El-Ansary, H.M. Abdel-Fattah, N.S. Abdel-Kader, Spectrochim. Acta A
79 (2011) 522–528.
[21] R. Amin, B. Krammer, N. Abdel-Kader, T. Verwanger, A. El-Ansary, Eur. J.
Medicinal Chem. 45 (2010) 372-378.
[22] A.I. Vogel, A.R. Tatchell, B.S. Furnis, A.J. Hannaford, P.W.G. Smith,
Vogel’s Textbook of Practical Organic Chemistry, fifth ed., Longmans, London
1989.
[23] A.M.G. Macdonald, P. Sirichanya, Microchem J. 14 (1969) 199-206.
[24] H.J. Emeleus, A.G. Sharpe, Modern Aspects of Inorganic Chemistry, fourth
ed., Roultedge and Paul, London, 1974.
[25] R.M. Amin, N.S. Abdel-Kader, A.L. El-Ansary, Spectrochim. Acta A 95
(2012) 517-525
[26] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81-122.
[27] P.G. Bradley, N. Kress, B.A. Hornberger, R.F. Dallinger, W.H. Woodruff, J.
Am. Chem. Soc. 103 (1981) 7441-7446.
[28] K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination
Compounds, Wiley, New York, 1986.
[29] N.S. Navaneetham, S. Soundararajan, Prec. Indian Acad. Sci. (Chem. Sci.) 89
(1980) 17-23.
[30] A.M. Ramadan, M.M. Ibrahim, Shaban, J. Mol. Struct. 1006 (2011) 348-355.
[31] S.B. Kalia, K. Lumba, G. Kaushal, M. Sharma, Indian J. Chem. 46 A (2007)
1233-1239.
[32] G. Speier, J. Csihony, A.M. Whalen, C.G. Pierpont, Inorg. Chem. 35 (1996)
3519-3524.
[33] D. Kivelson, R. Neiman, J. Chem. Phys. 35 (1961) 149-155.
[34] B.J. Hathaway, D.E. Billing, Coord. Chem. Rev. 5 (1970) 143-207.
[35] N.M. El-Metwally, I.M. Gabr, A.M. Shallaby, A.A. El-Asmy, J. Coord. Chem.
58 (2005) 1145-1159.
[36] A.W. Coats, J.P. Redfern, Nature 201 (1964) 68-69.
15

[37] H.H. Horowitz, G. Metzger, J. Anal. Chem. 35 (1963) 1464-1468.
[38] M. Nathg, P. Arora, Synth. React. Inorg. Met. Org. Chem. 23 (1993) 1523-
1545.
[39] T. Taakeyama, Fx. Quinn, Thermal Analysis, Fundamentals and Applications
to Polymer Science, Joh Wiley and Sons, Chichester, 1994.
[40] A.A. Frost, R.G. Pearson, Kinetics and Mechanisms, Wiley, New York, 1961.
[41] D. Greenwood, Antimicrobial chemothereapy, Bailliere, Tin dall, London.
Part II, Laboratory Aspects of Antimicrobial Therapy, 1983.
[42] V. Lorian, Antibiotics in Laboratory Medicine, Williams and Wilkins,
Baltimore, 1996.
16

1.2
1.0
0.8
0.6
0.4
0.2
0.0
H L
2
Cu-H L
2
Ni-H L
2
Co -H L
a
Co -H L
2
p
2
300
400
500
Wavelength (nm)
Fig. 1. Electronic absorption spectra of H2L ligands and its complexes
600
700
800
17

1.0
0.8
0.6
0.4
0.2
0.0
H4L
Cu-H4L
Ni-H4L
Coa-H4L
Cop-H4L
300
400
500
600
700
800
Wavelength (nm)
Fig. 2. Electronic absorption spectra of H4L ligands and its complexes
18

Fig. 3 TG plot ofCop-H2L, Coa-H2L, Coa-H4L and Nip-H4L complexes
19

Table 1: Analytical and physical data of the metal complexes of Schiff bases
(Calc.) Found
Molar
Ligand or complex
Symbol
Formula
conductance
ohm-1 cm2 mole-1
% C
% H
% N
4.95
% M
-
% Cl
66.32
4.40
H2L
[Cu(L)(H2O)2]2
H2L
C30H24O8N2
C60H52Cu2N4O20
C60H52Co2N4O20
C32H51CoN2O15Cl
C30H26NiN2O10
-
-
-
-
(66.66) (4.48) (5.18)
56.88 3.67 4.20
(56.47) (4.08) (4.39)
55.35 4.06 4.20
(56.88) (4.11) (4.43)
52.25 4.65 4.20
(52.63) (4.52) (3.84)
57.00 4.60 4.68
(56.99) (4.11) (4.43)
9.89
(9.97)
9.17
Cu-H2L
Coa-H2L
Cop-H2L
Ni-H2L
32.6
19.0
15.4
32.2
[Coa(L)]2.4H2O
(9.31)
7.86
2.71
[Cop(HL)(ClO4)½(C2H5OH)].2H2O
[Ni(HL)OH(H2O)]
(7.39)
9.07
(2.43)
-
(9.28)
The metal salt used in the preparation was: a = metal acetate, p = metal perchlorate

Table 1 continued
H4L
65.69
(65.62)
55.13
3.79
(3.93)
3.82
5.37
(5.47)
4.46
H4L
C28H20N2O8
C28H22CuN2O10
C56H48Co2N4O22
C28H26CoN2O12
C28H26NiN2O12
-
-
-
-
-
-
-
10.01
(10.42)
9.61
[Cu(H3L)OH(H2O)]
[Coa(H2L)]2.6H2O
Cu-H4L
Coa-H4L
Cop-H4L
Ni-H4L
34.8
11.7
19.7
22.0
(55.12)
53.80
(3.61)
3.85
(4.59)
4.90
(53.93)
53.40
(3.85)
4.15
(4.49)
4.50
(9.46)
9.00
[Cop(H3L)OH(H2O)].2H2O
[Ni(H3L)OH(H2O)3]
(52.42)
52.90
(4.06)
3.95
(4.37)
4.50
(9.19)
8.64
(52.25)
(4.06)
(4.37)
(9.16)

Table 2: Important IR bands of the ligands and their metal complexes
ν(H2O)
Compounds
ν(C=N)
ν(C=O)
ν(C-N)
ν(C-O)
ν(M-O)
ν(M-N)
ν (OH)
-
-
H2L
[Cu(L)(H2O)2]2
3433
3416
3383
3501
3380
3430
3404
3361
3379
3381
1602
-
1658
1658
1653
1658
1658
1656
1655
1663
1656
1657
1340
1315
1264
1302
1304
1341
1343
1344
1341
1342
1187
1235
1229
1226
1227
1170
1267
1222
1223
1267
559
558
547
549
-
484
486
473
473
-
[Coa(L)]2.4H2O
1611
-
[Cop(HL)(ClO4)½(C2H5OH)].2H2O
[Ni(HL)OH(H2O)]
H4L
-
1632
1627
1619
1620
1619
556
555
553
553
474
474
470
473
[Cu(H3L)OH(H2O)]
[Coa(H2L)]2.6H2O
[Cop(H3L)OH(H2O)].2H2O
[Ni(H3L)OH(H2O)3]

Table 3: Electronic absorption bands and magnetic moments of the Schiff bases
and their metal chelates
ν cm-1
BM
Compounds
H2L
30769, 25316
27027
-
Cu-H2L
Ni-H2L
Coa-H2L
Cop-H2L
H4L
2.21
0.00
6.36
2.49
-
25706
26882
25906
30769, 25510
26809
Cu-H4L
Ni-H4L
Coa-H4L
Cop-H4L
1.74
3.21
6.07
2.89
25706
25510
26109
23

Table 4: Electron spin resonance data of Cu (II) complexes
gꢀ
Complex
g||
gav
G
Cu-H2L
Cu-H4L
2.09
2.18
2.03
2.04
2.06
2.09
3.13
4.15
24

Table 5: Thermal behavior of some solid chelates with Schiff bases
Dehydration stage
Metal oxide
residue %
Decoordination stage
Complexes
Symbol M:L
Temp.
% weight loss
Calc. Found
Temp.
range
-
% weight loss
Calc. Found
Temp. °C Calc. Found
range
[Coa(L)]2.4H2O
Coa-H2L
2:2
1:1
2:2
1:1
50-130
5.69
4.93
8.67
-
5.56
5.00
8.33
-
-
-
570
595
596
597
11.84 11.66
10.26 10.00
12.03 12.22
11.17 11.00
[Cop(HL)(ClO4)½(C2H5OH)].2H2O Cop-H2L
50-130
50-180
-
130-330 6.30
6.11
-
[Coa(H2L)]2.6H2O
Coa-H4L
Ni-H4L
-
-
[Ni(H3L)OH(H2O)3]
150-300 8.07
8.00
25

Table 6: Kinetic data of the thermal decomposition of some H2L and H4L metal chelates
Correlation
Coefficient
r2
Decomposition
Temp. (K)
Ea
A
(S-1)
H*
S*
G*
Complex
Symbol
Cs
(KJmol-1)
(KJmol-1) (Jmol-1K-1) (KJmol-1)
Coa(L)]2.4H2O
Coa-H2L
493-673
573-853
513-698
613-723
122.31
71.79
72.77
83.11
5.03 E8
3.33 E6
3.53 E5
1.53 E5
116.77
67.07
67.09
77.28
-85.04
-125.40
-145.62
-152.77
173.40
138.18
166.55
184.36
0.36
0.37
0.37
0.35
0.98
[Cop(HL)(ClO4)½(C2H5OH)].2H2O Cop-H2L
0.98
0.99
0.99
[Coa(L)]2.6H2O
Coa-H4L
Ni-H4L
[Ni(H3L)OH(H2O)3]
26

Table 7: Biological activities of Cu(II) and Co(II) complexes of two Schiff bases
Inhibition zone diameter (mm/mg sample)
Sample
Escherichia Stphylococcus Aspergillus flavus Candida albicans
coli (G-)
aurreus (G+)
(Fungus)
0.0
(Fungus)
0.0
Control: DMSO
Tetracycline
(Antibacterial agent)
Amphotericin B
(Antifungal agent)
Cu-H2L
0.0
0.0
31.0
-
29.0
-
-
-
19.0
21.0
10.0
13.0
0.0
12.0
12.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
Coa-H2L
Cu-H4L
Coa-H4L
12.0
11.0
27

Research highlights
The structure of the prepared complexes was studied using many tools of
analysis.
The metal complexes exhibit different geometrical arrangements.
All formed complexes are 1:1 or 2:2 (M:L) and non-electrolyte chelates.
The tested complexes show antibacterial activity.
28

Benzopyran-4-one derivative
29