Metal complexes of Schiff base derived from sulphametrole and o-vanilin. Synthesis, spectral, thermal characterization and biological activity

Article abstract of DOI:10.1016/j.saa.2006.04.033

Metal complexes of Schiff base derived from condensation of o-vanilin (3-methoxysalicylaldehyde) and sulfametrole [N¹-(4-methoxy-1,2,5-thiadiazole-3-yl)sulfanilamide] (H₂L) are reported and characterized based on elemental analyses,

Full text of DOI:10.1016/j.saa.2006.04.033

Spectrochimica Acta Part A 66 (2007) 949–958
Metal complexes of Schiff base derived from sulphametrole and o-vanilin
Synthesis, spectral, thermal characterization and biological activity
Gehad G. Mohameda, Carmen M. Sharabyb,∗
a Chemistry Department, Faculty of Science, Cairo University, Giza, Egypt
b Chemistry Department, Faculty of Science, Al-Azhar University (Girls), Nasr City, Cairo, Egypt
Received 24 January 2006; accepted 17 April 2006
Abstract
Metal complexes of Schiff base derived from condensation of o-vanilin (3-methoxysalicylaldehyde) and sulfametrole [N1-(4-methoxy-1,2,5-
thiadiazole-3-yl)sulfanilamide] (H2L) are reported and characterized based on elemental analyses, IR, 1H NMR, solid reflectance, magnetic
moment, molar conductance, mass spectra, UV–vis and thermal analysis (TGA). From the elemental analyses data, the complexes were pro-
posed to have the general formulae [M2X3(HL)(H2O)5]·yH2O (where M = Mn(II), Co(II), Ni(II), Cu(II), Zn(II) and Cd(II), X = Cl, y = 0–3);
[Fe2Cl5(HL)(H2O)3]·2H2O; [(FeSO4)2(H2L)(H2O)4] and [(UO2)2(NO3)3(HL)(H2O)]·2H2O. The molar conductance data reveal that all the metal
chelates were non-electrolytes. The IR spectra show that, H2L is coordinated to the metal ions in a tetradentate manner with ON and NO donor sites
of the azomethine-N, phenolic-OH, enolic sulphonamide-OH and thiadiazole-N. From the magnetic and solid reflectance spectra, it is found that
the geometrical structures of these complexes are octahedral. The thermal behaviour of these chelates shows that the hydrated complexes losses
water molecules of hydration in the first step followed immediately by decomposition of the anions and ligand molecules in the subsequent steps.
The activation thermodynamic parameters, such as, E*, ꢀH*, ꢀS* and ꢀG* are calculated from the DrTG curves using Coats–Redfern method. The
synthesized ligand, in comparison to their metal complexes also were screened for their antibacterial activity against bacterial species, Escherichia
coli, Salmonella typhi, Bacillus subtillus, Staphylococcus aureus and Fungi (Aspergillus terreus and Aspergillus flavus). The activity data show
that the metal complexes to be more potent/antimicrobial than the parent Shciff base ligand against one or more microbial species.
© 2006 Elsevier B.V. All rights reserved.
Keywords: Sulphametrole; o-Vanilin; Transition metal complexes; IR; 1H NMR; Conductance; Solid reflectance; Magnetic moment; Thermal analysis; Biological
activity
1. Introduction
naphthaldehyde and benzoyl acetone with various sulfa drugs,
namely sulfadiazole, sulfadiazine, sulfisoxazole and sulfaguani-
dine have been prepared [2]. The stereochemical and biochem-
ical aspects of the complexes of sulfonamide imine of silicon
and tin have been studied [3].
Heterocyclic compounds containing both nitrogen and sul-
phur in the ring system have been utilized in synthesis of biolog-
ically active metal complexes. The biological activity of sulpha
drugs as well as the ligands drived from them has been known a
long time. Some of these drugs have now been observed to have
an increased biological activity when administered in the form
of metal complexes. The condensation products of sulpha drugs
with aldehydes, ketones or their derivatives are very active bio-
logically, besides having good complexing ability. Their activity,
too, increases on complexation with metal ions [1].
2. Experimental
2.1. Materials and reagents
All chemicals used were of the analytical reagent grade (AR),
and of highest purity available. They included sulphametrole
(Sigma); o-vanilin (Sigma); copper(II) chloride dihydrate (Pro-
labo); cobalt(II) and nickel(II) chloride hexahydrates (BDH);
zinc(II) and Cd(II) chloride dihydrates (Ubichem), uranyl nitrate
hexahydrate (Sigma) and ferric chloride hexahydrate (Prolabo).
Zinc oxide, disodium salt of ethylenediaminetetraacetic acid,
EDTA (Analar); ammonia solution (33% v/v) and ammonium
Theorganotin(IV)complexesoftheSchiffbasesderivedfrom
condensation of 2-acetylfuran, 2-acetylthiophene, 2-yhdroxy-1-
∗
Corresponding author.
E-mail address: csharaby@yahoo.com (C.M. Sharaby).
1386-1425/$ – see front matter © 2006 Elsevier B.V. All rights reserved.
doi:10.1016/j.saa.2006.04.033

950
G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
chloride (El-Nasr pharm. Chem. Co., Egypt). Organic solvents
used included absolute ethyl alcohol, diethylether, and dimethyl-
formamide (DMF). These solvents were either spectroscopic
pure from BDH or purified by the recommended methods [4]
and tested for their spectral purity. Hydrogen peroxide, sodium
chloride, sodium carbonate and sodium hydroxide (A.R.) were
used. Hydrochloric and nitric acids (MERCK) were used. De-
ionized water collected from all glass equipments was usually
used in all preparations.
2.2. Instruments
The spectrophotometric measurements in solution were car-
ried out using automated spectrophotometer UV–vis Perkin-
Elmer Model Lambda 20 ranged from 200 to 900 nm. The molar
conductance of solid complexes in DMF was measured using
Sybron–Barnstead conductometer (Meter-PM.6, E = 3406). Ele-
mental microanalyses of the separated solid chelates for C,
H, N and S were performed in the Microanalytical Center,
Cairo University. The analyses were repeated twice to check
the accuracy of the data. Infrared spectra were recorded on a
Perkin-Elmer FT-IR type 1650 spectrophotometer in wave num-
ber region 4000–200 cm−1. The spectra were recorded as KBr
discs. The solid reflectance spectra were measured on a Shi-
madzu 3101pc spectrophotometer. The molar magnetic suscep-
tibility was measured on powdered samples using the Faraday
method. The diamagnetic corrections were made by Pascal’s
constant and Hg[Co(SCN)4] was used as a calibrant. The ther-
mogravimetric analysis (TGA and DrTGA) was carried out in
dynamic nitrogen atmosphere (20 mL/min) with a heating rate of
10 ◦C min−1 using Shimadzu TGA-50H thermal analyzers. The
mass spectra were recorded by the EI technique at 70 eV using
MS-5988 GS–MS Hewlett-Packard instrument in the Microan-
1
alytical Center, Cairo University. The H NMR spectra were
recorded using 300 MHz Varian-Oxford Mercury.
2.3. Synthesis of Schiff base (H2L)
Hot solution (≈60 ◦C) of o-vanilin (7.6 g, 50 mmol) was
mixed with hot solution (≈60 ◦C) of sulphametrole (13.3 g,
50 mmol) in 50 mL ethanol. The resulting mixture was left under
reflux for 2 h and the formed solid product was separated by fil-
tration, purified by crystallization from ethanol, washed with
diethyl ether and dried in a vacuum over anhydrous calcium
chloride. The yellow product is produced in 88% yield (Table 1).
2.4. Synthesis of metal complexes
The metal complexes of the Schiff base, H2L were prepared
by the addition of hot solution (≈60 ◦C) of the appropriate metal
chloride or nitrate (2 mmol) in an ethanol–water mixture (1:1,
25 mL) to the hot solution (≈60 ◦C) of the Schiff base (0.42 g,
1 mmol) in the same solvent (25 mL). The resulting mixture was
stirred under reflux for one hour whereupon the complexes pre-
cipitated. They were collected by filtration, washed with a 1:1
ethanol:water mixture and diethyl ether. The analytical data for
C, H, N and S were repeated twice (Table 1).

G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
951
studies, molar conductance, mass spectra and thermal analy-
sis (TGA and DrTG), to identify their tentative formulae in
a trial to elucidate their molecular structures. The results of
elemental analyses, Table 1 are in good agreement with those
required by the proposed formulae. The formation of these com-
plexes may proceed according to the following equations given
below:
M2Cl3+H2L + 5H2O + yH2O→[M2Cl3(HL)(H2O)5]·yH2O,
M = Mn(II), Co(II), Ni(II), Cu(II), Zn(II), Cd(II); y = 0–3
Fig. 1. Schiff base (H2L) ligand.
2FeSO4 + H2L + 4H2O → [(FeSO4)2(H2L)(H2O)4]
2.5. Biological activity
2FeCl3 + H2L + 5H2O →[Fe2Cl5(HL)(H2O)3] · 2H2O
+ HCl
About 0.5 mL spore suspension (10−6 to 10−7 spore/mL) of
each of the investigated organisms was added to a sterile agar
medium just before solidification, then poured into sterile petri
dishes (9 cm in diameter) and left to solidify. Using sterile cork
borer (6 mm in diameter), three holes (wells) were made in each
dish, then 0.1 mL of the tested compounds dissolved in DMF
(100 (g/mL) were poured into these holes. Finally, the dishes
were incubated at 37 ◦C for 48 h where clear or inhibition zones
were detected around each hole. About 0.1 mL DMF alone was
used as a control under the same condition for each organism
and by subtracting the diameter of inhibition zone resulting with
DMF from that obtained in each case, both antimicrobial activ-
ities can be calculated as a mean of three replicates [5,6].
2UO2(NO3)2 + H2L + 3H2O → [(UO2)2(NO3)3(HL)(H2O)]
· 2H2O + HNO3
3.2.1. Molar conductivity measurements
The chelates were dissolved in DMF and the molar conduc-
tivities of 10−3 M of their solutions at 25 2 ◦C were measured.
Table 1, show the molar conductance values of the complexes. It
is concluded from the results that the chelates are found to have
molar conductance values of 7.78–16.32 ꢁ−1 mol−1 cm2 indi-
cating that these chelates are non-electrolytes. It also, indicates
the non-bonding of the chloride, sulfate or nitrate anions to the
metal ions.
3. Results and discussion
3.1. Schiff bases characterization
3.2.2. IR spectra and mode of bonding
The Schiff base, H2L, is subjected to elemental analyses.
The results of elemental analyses (C, H, N, S) with molecular
formula and the melting point are presented in Table 1. The
results obtained are in good agreement with those calculated for
the suggested formulae and the melting point is sharp indicating
the purity of the prepared Schiff base. The structure of the Schiff
base under study is given below in Fig. 1.
The structure of this Schiff base is also confirmed by IR and
1H NMR spectra, which will be discussed in detailed manner
with metal complexes latter. The electron impact mass spec-
trum of H2L was recorded and investigated at 70 eV of electron
energy. The mass spectrum of H2L shows a well-defined par-
ent peak at m/z = 420 with a relative intensity = 5%. The parent
ion and the fragments at m/z (%): 297 (5), 290 (10), 226 (10)
and 156 (100) were obtained by cleavage in different positions in
H2L molecule, are shown in Scheme 1, confirming the suggested
structure of the H2L ligand.
The IR data of the spectra of H2L Schiff base and its com-
plexes are listed in Table 2. The IR spectra of the complexes
are compared with those of the free ligand in order to deter-
mine the coordination sites that may involve in chelation. There
are some guide peaks, in the spectra of the ligand, which are
of good help for achieving this goal. The position and/or the
intensities of these peaks are expected to be changed upon
chelation. New peaks are also guide peaks as well as water in
chelation. These guide peaks are listed in Table 2. Upon com-
parison it was found that the azomethine υ(C N) stretching
vibration is found in the free ligand at 1632 cm−1. This band
is shifted to lower wavenumbers (1591–1609 cm−1) in the com-
plexes indicating the participation of the azomethine nitrogen
in coordination (M N) [7]. Phenolic υ(C O) stretching vibra-
tion band is observed at 1306 cm−1 in the free ligand. In all
complexes, this band appears at higher (or lower) frequencies
in 1046–1100 cm−1 region, confirming the involvement of the
phenolic group in complex formation [8].
3.2. Composition and structures of Schiff base complexes
The stretching vibration band; ν(NH), of the sulfonamide
group, which found at 3412 cm−1 in the free ligand, was dis-
appeared or hidden under the peak of water molecules. The
presence of coordinated water molecules renders it difficult to
confirm the enolization of the sulfonamide group. The SO2
group modes of the ligand appear as sharp bands at 1328 and
The isolated solid complexes of Fe(III), Mn(II), Fe(II),
Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and UO2(II) ions with
the Schiff base H2L ligand were subjected to elemental anal-
1
yses (C, H, N, S and metal content), IR, H NMR, magnetic

952
G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
Scheme 1. The possible fragmentation pathways of H2L ligand.
1155 cm−1 (νasym(SO2)) and (νsym(SO2)), respectively. In the
complexes, the asymmetric and symmetric modes are shifted
to 1312–1322 and 1115–1137 cm−1, respectively, upon coordi-
nation to the transition metals [6–9]. The blue shift of the SO2
stretching vibration to lower frequencies may be attributed to the
transformation of the sulfonamide ( SO2NH) to give the enol
form ( SO(OH) N) as a result of complex formation to give
more stable six-membered ring [9–12].
the metal complexes. These shifts refer to the coordination
through thiadiazole-N atom. The IR bands at 812–850 and
780–796 cm−1, υ(H2O) of coordinated water, is an indication
of the binding of the water molecules to the metal ions. New
bands are found in the spectra of the complexes in the regions
520–586, which are assigned to υ(M O) stretching vibrations.
The bands at 434–466 have been assigned to υ(M N) mode
[13].
A sharp band due to the υ(C N) stretching vibration
of thiadiazole ring appeared at 1570 cm−1 in H2L ligand.
This band is shifted to 1231–1274 cm−1 in the spectra of
Therefore, fromtheIRspectra, itisconcludedthatH2Lligand
behaves as a uni-negtively tetradentate ligand coordinated to the
metal ions via the azomethine-N, deprotonated phenolic-O, eno-

G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
953
Table 2
IR spectra (4000–400 cm−1) of the H2L ligand and its metal complexes
Compound
ν(NH) ν(OH)
ν(SO2)
ν(SO2)
ν(CH N) ν(C O)
ν(C N)
ν(H2O)
ν(M O)
ν(M N)
enolic
(asym)
(sym.)
azome-
thine
phenolic
thiodiaza
(coord.)
H2L
3412br
–
–
1328sh
1313m
1155sh
1126m
1632sh
1591br
1306sh
10.92m
1156sh
1158sh
–
–
–
3040w
3046w
–
844m
796sh
583m
450w
[Fe2Cl5(HL)(H2O)3]·2H2O
–
–
–
–
–
–
–
–
1315m
1322w
1313m
1321m
1312m
1320w
1320m
1322m
1143s
11598m
1595sh
1596sh
1609br
1595m
1596sh
1594m
1592m
1060s
1170w
1156br
1188br
1188br
1156sh
1168sh
1168br
1168br
820s
782s
554w
521m
520m
550m
586m
520m
545m
547m
440w
436w
434w
450w
466w
461w
462w
462w
[Mn2Cl3(HL)(H2O)5]
[(FeSO4)2(H2L)(H2O)4]
[Co2Cl3(HL)(H2O)5]·2H2O
[Ni2Cl3(HL)(H2O)5]
1120w
1115w
1130w
1125m
1137w
1125m
1125m
1090s
826s
790s
3045w
3050w
3041w
3068w
3070w
3031w
1052m
1046m
1092sh
1080sh
1100s
850m
780m
820m
780m
842m
796sh
[Cu2Cl3(HL)(H2O)5]·3H2O
[Zn2Cl3(HL)(H2O)5]
812m
782m
840m
795m
[Cd2Cl3(HL)(H2O)5]
[(UO2)2(HL)(NO3)3(H2O)]·2H2O
1086m
–
sh, sharp; m, medium; s, small; w, weak; br, broad.
lic sulphonamide-OH and thiadiazole-N; except Fe(II) complex,
it act as neutral and coordinated via protonated phenolic-OH.
Table 3
Characteristic 1H NMR spectra for the H2L and its Zn(II) and Cd(II) metal
complexes
3.2.3. 1H NMR spectra
Compound
Chemical
Assignment
A survey of literature revealed that the 1H NMR spectroscopy
has been proved useful in establishing the nature and structure of
many Schiff bases as well as their complexes in solutions. The
1H NMR spectrum of Schiff base was recorded in dimethylsul-
foxide (DMSO-d6) solution using tetramethylsilane (TMS) as
internal standard. The 1H NMR spectra of the Schiff base H2L,
its diamagnetic Zn(II) and Cd(II) complexes and the chemical
shifts of the different types of protons are listed in Table 3. The
spectra of the complexes are examined in comparison with those
of the parent Schiff base. Upon examinations it was found that
the phenolic-OH signal, appeared in the spectrum of H2L lig-
and at 12.82 ppm is completely disappeared in the spectra of
its Zn(II) and Cd(II) complexes indicating that the OH proton
is removed by the chelation with Zn(II) and Cd(II) ions. The
signal observed at 3.33 and 3.75 ppm with an integration corre-
sponding to ten protons in case of Zn(II) and Cd(II) complexes,
respectively, are assigned to five water molecules.
shift, δ (ppm)
H2L
4.00
6.58
6.95
7.43
8.70
9.77
11.04
s, 6H, 2 OCH3
d, 2H, ArH’s, J = 8.70 Hz
m, 2H, vanillin-H
m, 1H, vanillin-H
d, 2H, ArH’s, J = 8.70 Hz
s, 1H, CH
N
brs; 1H, NH (exchangeable with
D2O)
s, 1H, OH-vanilin (exchangeable
with D2O)
12.82
[Zn2Cl3(HL)(H2O)5]
3.75
4.00
6.00
s, coordinated water
s, 6H, 2 OCH3
br, 1H, OH-enolic (exchangeable
with D2O)
d, 2H, ArH’s, J = 8.85 Hz
m, 2H, vanillin-H
m, 1H, vanillin-H
6.57
6.98
7.46
7.57
9.78
d, 2H, ArH’s, J = 8.85 Hz
s, 1H, CH
N
[Cd2Cl3(HL)(H2O)5]
3.33
4.00
6.00
s, coordinated water
s, 6H, 2 OCH3
brs, 1H, OH-enolic
(exchangeable with D2O)
d, 2H, ArH’s, J = 8.70 Hz
m, 2H, vanillin-H
3.2.4. Magnetic susceptibility and electronic spectra
measurements
The UV–vis spectra of the ligand and the complexes were
recorded in DMF solution in the wavelength range from 200 to
800 nm. The spectra showed a sharp and intense band at 207 nm
6.56
6.97
7.45
7.56
9.78
*
m, 1H, vanillin-H
d, 2H, ArH’s, J = 8.70 Hz
which is attributed to ␲–␲ transition within the benzene ring
of the ligand molecule. This band is red shifted to 221–237 nm.
The two bands observed at 302 and 345 nm in the free ligand
s, 1H, CH
N

954
G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
are reasonably accounted for ␲–␲* and n–␲* transitions for the
phenolic-OH and azomethine moieties, respectively. The blue or
red shifts of these bands in the regions 295–305 and 331–365 nm
with respect to the ligand depend on the type of metal ions coor-
dinated to the ligand and indicated coordination of phenolic-OH
and azomethine moieties to the metal ions. The spectra of the
complexes further display a bands in the range 413–495 and
559–642 nm, which might be assigned to charge transfer transi-
tion from the ligand to metal ions (L → M) [12].
try [17]. The electronic spectrum of the Co(II) complex give
three bands at 15,576, 17,467 and 21,930 cm−1 wave num-
ber regions. The bands observed are assigned to the tran-
4
4
4
4
sitions T1g (F) → T2g(F) (ν1), T1g(F) → A2g(F) (ν2) and
4
4T1g(F) → T2g(P) (ν3), respectively, suggesting that there is an
octahedral geometry around Co(II) ion [18–20]. The magnetic
susceptibility measurements lie at 5.20 B.M. (normal range for
octahedral Co(II) complexes is 4.3–5.2 B.M.), is an indicative
of octahedral geometry [21]. The region at 24,390 cm−1 refers
to the charge transfer band.
From the diffuse reflectance spectrum it is observed that, the
Fe(III) chelate exhibit a band at 20,942 cm−1, which may be
assigned to the 6A1g → T2g(G)transitionin octahedralgeometry
The Ni(II) complex reported herein has a room temperature
magnetic moment value of 3.67 B.M.; which is in the normal
range observed for octahedral Ni(II) complexes (μeff = 2.9–3.3
B.M.) [17]. This indicates that, the complexes of Ni(II) are
six coordinate and probably octahedral [22]. The electronic
5
of the complexes [14]. The 6A1g → T1g transition appears to be
split into two bands at 17,467 and 15,174 cm−1. The observed
magnetic moment of Fe(III) complex is 5.62 B.M. Thus, the
complexes formed have the octahedral geometry [15]. The spec-
trum shows also a band at 28,531 cm−1 which may attribute to
ligand to metal charge transfer. The diffuse reflectance spec-
trum of the Mn(II) complex shows three bands at 15,873, 21,739
*
*
spectrum, in addition to show the ␲–␲ and n–␲ bands of
free ligands, displays three bands, in the solid reflectance
3
3
spectrum at ν1: 15,698 cm−1: A2g → T2g; ν2: 17,422 cm−1
:
3A2g → T1g(F) and ν3: 20,202 cm−1
:
3A2g → T1g(P). The
3
3
spectrum shows also a band at 23,255 cm-1 which may attribute
to ligand to metal charge transfer. The reflectance spectrum of
Cu(II) chelat consists of a broad, low intensity shoulder band
centered at 16,537 and 17,467–22,422 cm−1. The 2Eg and 2T2g
states of the octahedral Cu(II) ion (d9) split under the influ-
ence of the tetragonal distortion and the distortion can be such
6
6
and 27,77 cm-1 assignable to 4T1g → A1g, 4T2g(G) → A1g and
4T1g(D) → A1g transitions, respectively [16]. The magnetic
6
moment value is 5.24 B.M. which indicates the presence of
Mn(II) complex in octahedral structure.
The electronic spectrum of the Fe(II) complex displays two
absorption bands at 15,480 and 17,376 cm−1 which are assigned
as to cause the three transitions 2B1g → B2g; 2B1g → Eg and
2
2
5
5
to T2g → Eg transitions [17]. Also the band at 23,255 cm−1
is assigned to L → M charge transfer. The observed magnetic
moment of 5.39 B.M. is consistent with an octahedral geome-
2B1g → A1g to remain unresolved in the spectra [23]. It is
2
concluded that, all three transitions lie within the single broad
Table 4
Thermogravimetric data of H2L metal complexes
Complex
TG range
(◦C)
DTGmax (◦C)
na
% Found (calcd.)
Mass loss
Assignment
Metallic
residue
Total mass
loss
30–100
100–250
55
160
1
1
4.54(4.50)
29.16 (28.96
Loss of 2H2O
Loss of 5HCl,
1/2H2, 1.5O2
Loss of HL
79.58
(79.99)
[Fe2Cl5(HL)(H2O)3]·2H2O
Fe2O3
2FeSO4
2CoO
250–700
280, 500
2
45.88 (46.53)
30–250
250–1000
70, 170
270, 670
2
2
8.92(9.03)
52.60 (52.82)
61.52
(61.85)
Loss of 4HCl
Loss of H2L
[(FeSO4)2(H2L)(H2O)4]
30–100
100–230
60
180
1
1
4.16(4.67)
16.59 (16.42)
Loss of 2H2O
Loss of HCl and
5H2O
80.84
(80.66)
[Co2Cl3(HL)(H2O)5] ·2H2O
230–650
30–140
270, 750
120
2
1
60.09 (59.57)
26.68(27.03)
Loss of HL
67.82
(67.60)
Loss of 3HCl and
C2H3NOS
[Ni2Cl3(HL)(H2O)5]
140–800
270, 590
2
41.14 (40.57)
Loss of 3H2O,
1.5H2 and HL
Loss of 3H2O
Loss of 3HCl,
3H2O, O2 and
1/2H2
2NiCO3
Cu2O
30–130
130–280
90
240
1
1
6.95(6.77)
36.14 (35.80)
83.24
(81.94)
[Cu2Cl3(HL)(H2O)5] ·3H2O
280–750
400, 650
2
40.15 (39.37)
Loss of HL
30–130
130–1000
50
1
3
2.82(3.00)
31.80 (31.33)
34.62
(34.33)
Loss of 2H2O
Loss of H2O and
HL
[(UO2)2(HL)(NO3)3–(H2O)]·2H2O
220, 400, 650
2UO2(NO3)2
a
n, number of decomposition steps.

G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
955
envelope centered at the same range previously mentioned. The
magnetic moment of 2.10 B.M. falls within the range normally
observed for octahedral Cu(II) complexes [24]. A moderately
intense peak observed at 23,364 cm−1 is due to ligand–metal
charge transfer transition [25].
ThecomplexesofZn(II), Cd(II)andUO2(II)arediamagnetic.
In analogy with those described for Zn(II) complexes containing
N–O donor Schiff bases [26–28] and according to the empirical
formulae of these complexes, we proposed an octahedral geom-
etry for these three complexes.
the removal of the organic part of the ligand leaving metal oxide
as a residue. The overall weight loss amounts to 79.58% (calcd.
79.99%).
The TGA curves of the Co(II) and Cu(II) chelates show four
stages of decomposition within the temperature range of 30–650
and 30–750 ◦C, respectively. The first stage at 30–100 ◦C cor-
responds to the loss of water molecules of hydration. The
energy of activation for this dehydration step was 39.02 and
63.85 kJ mol−1 for Co(II) and Cu(II) complexes, respectively.
While the subsequent (second, third and fourth) stages involve
the loss of 1–3HCl, 3–5H2O, 1/2H2, O2 and ligand molecules
with a mass loss of 76.68% (calcd. 75.99%) and 76.29% (calcd.
75.17%) for Co(II) and Cu(II) complexes, respectively. The
overall weight loss amounts to 80.84% (calcd. 80.66%) and
83.24% (calcd. 81.94%) for Co(II) and Cu(II) chelates, respec-
tively.
On the other hand [(FeSO4)2(H2L)(H2O)4] chelate exhibits
four decomposition steps. The first step in the temperature
range 30–250 ◦C (mass loss = 8.92%; calcd. for 4H2O; 9.03%)
may accounted for the loss of water molecules of coordi-
nation. The energy of activation for these steps was 53.25
and 82.37 kJ mol−1 for the first and second steps, respec-
tively. As shown in Table 4, the mass losses of the remain-
ing decomposition steps amount to 52.60% (calcd. 52.82%).
They correspond to the removal of H2L molecules leaving
FeSO4 as a residue. The energy of activation for these steps
was 132.4 and 176.9 kJ mol−1 for the third and fourth steps,
respectively.
4. Thermal analyses (TGA and DrTG)
In the present investigation, heating rates were suitably
controlled at 10 ◦C min−1 under nitrogen atmosphere and the
weight ◦loss was measured from the ambient temperature up to
∼
1000 C. The data are listed in Table 4. The weight losses for
=
each chelate were calculated within the corresponding tempera-
ture ranges. The different thermodynamic parameters are listed
in Table 5.
The thermogram of [Fe2Cl5(HL)(H2O)3]·2H2O chelate
shows four decomposition steps within the temperature range
30–700 ◦C. The first two steps of decomposition within the
temperature range 30–250 ◦C correspond to the loss of water
molecules of hydration, 5HCl, 1/2H2 and 1.5O2 gases with a
mass loss of 33.70% (calcd. 33.46%). The energy of activation
was 44.40 and 78.24 kJ mol−1 for the first and second steps,
respectively. The subsequent steps (250–700 ◦C) correspond to
Table 5
Thermodynamic data of the thermal decomposition of H2L metal complexes
Complex
Decomp.
temperature
(◦C)
E*
(kJ mol−1
A (s−1
)
ꢀS*
(J K−1 mol−1
ꢀH*
(kJ mol−1
ꢀG*
(kJ mol−1
)
)
)
)
30–100
130–230
230–320
400–550
44.40
78.24
119.6
182.7
3.67 × 108
4.09 × 1012
5.89 × 1010
2.99 × 109
−49.15
−85.76
−112.3
−165.9
62.18
59.86
122.3
196.8
64.57
88.98
142.1
205.7
[Fe2Cl5(HL)(H2O)3]·2H2O
30–110
140–230
240–330
550–700
53.25
82.37
132.4
176.9
2.87 × 108
4.55 × 1012
6.08 × 105
9.34 × 107
−34.65
−62.66
−106.6
−185.7
45.65
97.16
125.3
161.5
37.22
76.42
109.5
183.4
[(FeSO4)2(H2L)–(H2O)4]
30–100
110–240
240–360
510–620
39.02
67.55
134.3
205.1
8.23 × 109
5.68 × 1010
6.75 × 107
6.03 × 1012
5.09 × 109
4.68 × 106
8.44 × 1010
−34.89
−76.55
−116.5
−177.5
−63.88
−108.2
−162.7
43.52
85.42
128.9
197.7
38.96
117.0
155.6
99.45
[Co2Cl3(HL)(H2O)5] ·2H2O
[Ni2Cl3(HL)(H2O)5]
30–140
200–310
450–620
22.95
75.98
137.4
60.52
96.74
166.3
48.89
141.2
186.9
30–110
150–270
300–460
500–720
63.85
145.6
187.9
165.7
6.53 × 1012
3.76 × 1010
7.02 × 108
1.98 × 109
−43.76
−89.77
−152.4
−168.7
67.32
106.1
166.8
225.4
68.66
111.3
152.6
183.7
[Cu2Cl3(HL)(H2O)5] ·3H2O
30–110
120–280
300–520
560–720
51.18
102.4
182.9
243.4
4.45 × 107
3.19 × 1012
4.56 × 1011
7.21 × 107
−56.14
−95.28
−166.7
−203.2
50.46
99.43
128.4
195.8
50.66
91.86
152.3
192.6
[(UO2)2(HL)(NO3)3–(H2O)]·2H2O

956
G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
[Ni2Cl3(HL)(H2O)5] complex was thermally decomposed
vation energies were 75.98 and 137.4 kJ mol−1 for the second
and third steps, respectively.
in three decomposition steps within the temperature range of
30–800 ◦C. The first decomposition step with an estimated mass
loss of 26.68% (calcd. mass loss = 27.03%) within the tempera-
ture range 30–140 ◦C, may be attributed to the liberation of the
The TGA curve of the UO2(II) chelate represent four decom-
position steps as shown in Table 4. The first step of decom-
position within the temperature range 30–130 ◦C corresponds
to the loss of hydrated water molecules with a mass loss of
2.82% (calcd. for 2H2O; 3.00%). The energy of activation for
this dehydration step was 51.18 kJ mol−1. The remaining steps
of decomposition within the temperature range 130–1000 ◦C
correspond to the removal of H2L ligand as gases with an
3HCl and C2H3NOS. The activation energy was 22.95 kJ mol−1
.
The remaining decomposition steps (two steps) found within
the temperature range 140–800 ◦C with an estimated mass loss
of 41.14% (calcd. mass loss = 40.57%) which are reasonably
accounted for by the removal of H2L ligand as gases. The acti-
Fig. 2. The proposed structures of H2L complexes.

G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
957
ꢀH∗ = E∗ − RT
(2)
(3)
energy of activation for these steps of 102.4, 182.9 and
243.4 kJ mol−1 for the second, third and fourth steps, respec-
tively. The overall weight losses amount to 31.80% (calcd.
31.33%).
ꢀG∗ = ꢀH∗ − TꢀS∗
The data are summarized in Table 5. The activation
energies of decomposition were found to be in the range
22.95–243.4 kJ mol−1. The high values of the activation ener-
gies reflect the thermal stability of the complexes. The entropy
of activation was found to have negative values in all the com-
plexes which indicates that the decomposition reactions proceed
with a lower rate than the normal ones.
5. Kinetic data
The thermodynamic activation parameters of decomposition
processes of dehydrated complexes namely activation energy
(E*), enthalpy (ꢀH*), entropy (ꢀS*) and Gibbs free energy
change of the decomposition (ꢀG*) were evaluated graphi-
callybyemployingtheCoats–Redfernrelation [29]. Theentropy
of activation (ꢀS*), enthalpy of activation (ꢀH*) and the free
energy change of activation (ꢀG*) were calculated using the
following equations:
6. Structural interpretation
The structures of the complexes of Schiff base with Fe(III),
Mn(II), Fe(II), Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and UO2(II)
ions were confirmed by the elemental analyses, IR, NMR, molar
conductance, magnetic, solid reflectance, UV–vis, mass and
thermal analysis data. Therefore, from the IR spectra, it is con-
ꢀ
ꢁ
ꢂꢃ
Ah
kT
ꢀS∗ = 2.303 log
R
(1)
Table 6
Antimicrobial activity of H2L and corresponding metal complexes
Samples
Concentration
(mg/mL)
Test organisms
Escherichia
coli
Salmonella
typhi
Staphylococcus
aureus
Bacillus
subtillus
Aspergillus
terreus
Aspergillus
flavus
1
2.5
5
++
++
+++
++
+++
+++
+
++
++
++
++
++
+
+
+
+
+
+
H2L
1
2.5
5
+
++
++
++
+++
+++
+
++
++
++
++
++
0
0
+
+
+
+
[Fe2Cl5(HL)(H2O)3]·2H2O
[(FeSO4)2(H2L)(H2O)4]
[Co2Cl3(HL)(H2O)5]·2H2O
[Ni2Cl3(HL)(H2O)5]
[Cu2Cl3(HL)(H2O)5]·3H2O
[Zn2Cl3(HL)(H2O)5]
[Cd2Cl3(HL)(H2O)5]
1
2.5
5
+
+
+
+
+
++
+
+
++
+
+
++
+
+
+
+
+
+
1
2.5
5
+
++
++
+
++
+++
+
+
++
++
++
+++
+
++
++
+
++
++
1
2.5
5
+
+
+
+
+
++
++
++
++
+
+
+
+
+
+
+
+
+
1
2.5
5
+
++
++
++
+++
+++
+
+
+
+
+
+
+
+
+
+
+
+
1
2.5
5
++
+++
+++
+++
+++
+++
+
+
++
++
++
++
+
+
+
0
0
+
1
2.5
5
++
+++
+++
+++
+++
+++
+
+
++
++
++
++
+
+
++
+
+
++
1
++
++
+++
+
+
++
+
+
+
+
+
+
0
+
+
+
+
+
[(UO2)2(HL)(NO3)3(H2O)]·2H2O 2.5
5
1
2.5
5
++
+++
+++
+++
+++
+++
++
++
+++
++
+++
+++
+++
+++
+++
+++
+++
+++
St.
St., reference standard; chloramphenicol was used as a standard antibacterial agent and grisofluvine was used as a standard antifungal agent. The test was done using
the diffusion agar technique. Inhibition values = 0.1–0.5 cm beyond control = +; inhibition values = 0.6–1.0 cm beyond control = ++; inhibition values = 1.1–1.5 cm
beyond control = +++; inhibition values > 1.5 cm beyond control = ++++; 0 = not detected. Well diameter 1 cm (100 ␮l of each conc. was tested).

958
G.G. Mohamed, C.M. Sharaby / Spectrochimica Acta Part A 66 (2007) 949–958
cluded that H2L behaves as a uninegatively tetradentate ligand
coordinated to the metal ions via the azomethine-N, deproto-
nated phenolic-OH (except Fe(II), protonated phenolic-OH),
enolic sulphonamide-OH and thiadiazole-N. The molar conduc-
tance data, it is found that the complexes are non-electrolytes.
[3] P. Nagpal, R.V. Singh, Appl. Organometal. Chem. (2004) 221.
[4] A.I. Vogel, Quantitative Inorganic Analysis Including Elemental Instru-
mental Analysis, 2nd ed., Longmans, London, 1962.
[5] M.E. Ibrahim, A.A.H. Ali, F.M.M. Maher, J. Chem. Technol. Biotechnol.
55 (1992) 217.
[6] N. Sari, S. Arslan, E. Logoglu, I. Sakiyan, J. Sci. 16 (2003) 283.
[7] G.G. Mohamed, Z.H. Abd El-Wahwb, J. Thermal Anal. 73 (2003)
347.
1
The H NMR spectra of the free ligand and its diamagnetic
Zn(II)complexshowthattheOHsignal, appearedinthespecrum
of H2L ligand at 3.34 ppm is completely disappeared in the
spectrum of its Zn(II) complex indicating that the OH pro-
ton is removed by the chelation with Zn(II) ion. On the basis
of the above observations and from the magnetic and solid
reflectance measurements, octahedral geometries are suggested
for the investigated complexes. As a general conclusion, the
investigated Schiff base behaves as a tetradentate and its metal
complexes structurs can be given as shown below (Fig. 2).
[8] N. Raman, A. Kulandaisamy, K. Jeyasubramanian, Synth. React. Inorg.
Met. Org. Chem. 31 (7) (2001) 1249.
[9] A.S. Salameh, H.A. Tayim, Polyhedron 2 (8) (1983) 829.
[10] H.A. Tayim, A.S. Salameh, Polyhedron 2 (10) (1983) 1091.
[11] B.T. Thaker, Proc. Natl. Acad. Sci. India, Sect. A 58 (3) (1988) 443.
[12] S.D. Kolwalkar, B.H. Mehta, Asian J. Chem. 8 (3) (1996) 406.
[13] J.R. Ferraro, Low Frequency Vibrations of Inorganic and Coordination
Compounds, 1st ed., Plenum Press, New York, 1971, p. 168.
[14] M.M. Moustafa, J. Thermal Anal. 50 (1997) 463.
[15] G.G. Mohamed, Nadia E.A. El-Gamel, F.A. Nour El-Dien, Synth. React.
Inorg. Met. Org. Chem. 31 (2) (2001) 347.
[16] S. Zhou, S. Liu, G. Zhou, Huaxue Shiji 23 (1) (2001) 26.
[17] F.A. Cotton, G. Wilkinson, C.A. Murillo, M. Bochmann, Advanced Inor-
ganic Chemistry, 6th ed., Wiley, New York, 1999.
7. Biological activity
In testing the antibacterial activity of these compounds we
used more than one test organism to increase the chance of
detecting antibiotic principles in tested materials. The sensitivity
of a microorganism to antibiotics and other antimicrobial agents
was determined by the assay plates which incubated at 28 ◦C
for 2 days for yeasts and at 37 ◦C for 1 day for bacteria. The
data showed that in some cases the ligand has a higher or similar
antmicrobial activity than the selected standards (chlorampheni-
col and grisofluvine). Also, the presence of certain metal ions
inhanced the antimicrobial activity of the ligand and in some
cases a higher or similar activity than the selected standards
were also showed (Table 6).
[18] G.G. Mohamed, N.E.A. El-Gamel, F. Tiexidor, Polyhedron 20 (2001) 2689.
[19] M.S. Masoud, A.M. Hindawy, A.S. Soayed, Trans. Met. Chem. 16 (1991)
372.
[20] M.M. Omar, G.G. Mohamed, Spectrochim. Acta (Part A) 61 (5) (2005)
929.
[21] N. Mondal, D.K. Dey, S. Mitra, K.M. Abdul Malik, Polyhedron 19 (2000)
2707.
[22] D.R. Zhu, Y. Song, Y. Xu, Y. Zhang, S.S.S. Raj, H.K. Fun, X.Z. You,
Polyhedron 19 (2000) 2019.
[23] J. Kohout, M. Hvastijova, J. Kozisek, L.G. Diaz, M. Valko, L. Jager, I.
Svoboda, Inorg. Chem. Acta 287 (1999) 186.
[24] N.R.S. Kumar, M. Nethiji, K.C. Patil, Polyhedron 10 (3) (1991) 365.
[25] J. Manonmani, R. Thirumurugan, M. Kandaswamy, M. Kuppayee, S.S.S.
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[26] J. Sanmartin, M.R. Bermejo, A.M.G. Deibe, M. Maneiro, C. Lage, A.J.C.
Filho, Polyhedron 19 (19) (2000) 185.
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