Synthesis, spectroscopic, electrochemical characterization, density functional theory (DFT), time dependent density functional theory (TD-DFT), and antibacterial studies of some Co(II), Ni(II), and Cu(II) chelates of (E)-4-(1,5-dimethyl-3-oxo-2-phenyl-2,3

Article abstract of DOI:10.1002/aoc.5729

In this study, mononuclear Co(II), Ni(II), and Cu(II) chelates of the potentially tridentate O, N, S donor ligand (E)-4-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-1-(3-hydroxynaphthalen-2-yl)methylene) thiosemicarbazide (H₂L) wer

Full text of DOI:10.1002/aoc.5729

Received: 11 February 2020
DOI: 10.1002/aoc.5729
Revised: 22 March 2020
Accepted: 28 March 2020
F U L L P A P E R
Synthesis, spectroscopic, electrochemical characterization,
density functional theory (DFT), time dependent density
functional theory (TD-DFT), and antibacterial studies of
some Co(II), Ni(II), and Cu(II) chelates of (E)-
4-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-
1-(3-hydroxynaphthalen-2-yl)methylene) thiosemicarbazide
Schiff base ligand
Ayman K. El-Sawaf1,2
| Yasser K. Abdel-Monem2 | Maged A. Azzam1,2
1Department of Chemistry, College of
Science and Humanities in Al-Kharj,
Prince Sattam Bin Abdulaziz University,
Al-Kharj, 11942, Saudi Arabia
2Department of Chemistry, Faculty of
Science, Menoufia University, Shebin El-
Kom, Egypt
Abstract
In this study, mononuclear Co(II), Ni(II), and Cu(II) chelates of the potentially
tridentate O, N, S donor ligand (E)-4-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-
1H-pyrazol-4-yl)-1-(3-hydroxynaphthalen-2-yl)methylene)
thiosemicarbazide
(H2L) were synthesized and characterized by elemental analyses, conductivity
measurements, Fourier transform infrared spectroscopy, UV–visible spectros-
copy, magnetic susceptibility, electron spin resonance (ESR) spectral analysis,
thermal thermogravimetric analyses, cyclic voltammetry, and theoretical den-
sity functional theory (DFT) and time dependent density functional theory
(TD-DFT) studies. The potential donors O (phenolic) and N (azomethine) are
found in a syn configuration due to a strong intramolecular hydrogen bond
[O–H …. N(1)] as shown in the 1H NMR spectrum of the free ligand. In
addition, analytical results revealed the formation of the metal chelates in a
ligand-to-metal molar ratio of 1:1. The organic ligand (H2L) coordinated to
the metal center through azomethine nitrogen, phenolate oxygen, and
thione/thiolate sulfur groups with the formation of a five- and six-membered
rings. The magnetic and spectral results supported the formation of tetrahe-
dral geometry around the Co(II) and Ni(II) centers, while Cu(II) centers were
assumed to have a square planar geometry. The time dependent density
functional theory (TD-DFT) calculation outputs were used to expect and
explain the experimental bathochromic and hypsochromic shifts resulting
from the ligand chelation. The antibacterial activity of the compounds under
study was evaluated in vitro against Staphylococcus aureus and Escherichia
coli demonstrating Gram-positive and Gram-negative bacteria, respectively. It
was found that among synthesized compounds, Cu(II) chelate 5 showed the
most inhibitory activity against both Gram-positive and Gram-negative
Correspondence
Ayman K. El-Sawaf, Department of
Chemistry, College of Science and
Humanities in Al-Kharj, Prince Sattam
Bin Abdulaziz University, Al-Kharj 11942,
Saudi Arabia.
Email: elsawaf2008@yahoo.com
Appl Organomet Chem. 2020;e5729.
wileyonlinelibrary.com/journal/aoc
© 2020 John Wiley & Sons, Ltd.
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https://doi.org/10.1002/aoc.5729

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EL-SAWAF ET AL.
bacteria when compared to the organic ligand as well as to the other pre-
pared chelates.
K E Y W O R D S
antibacterial activity, metal chelates, TD-DFT, thermal studies, Thiosemicarbazone
1 | INTRODUCTION
and its increased (decreased) as described for many tran-
sition metal chelates.[7–9] In continuation of our ongoing
During the past five decades, there have been much
development of research in the coordination chemistry of
nitrogen–oxygen–sulfur donor ligands such as substituted
interest in the synthesis of new thiosemicarbazone
chelates,[10,11] we report here the synthesis, structural
characterization, and antibacterial studies of new Co(II),
Ni(II), and Cu(II) chelates of (E)-4-(1,5-dimethyl-3-oxo-
2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-1-(3-hydro-
xynaphthalen-2-yl)methylene) thiosemicarbazide Schiff
base ligand (Scheme 1).
thiosemicarbazides,
thiosemicarbazones,
and
dithiocarbazates. Thiosemicarbazones, especially aro-
matic and heteroaromatic componds, have been used in
several applications as depending on the ligand donor
atoms, their flexibility and capability to coordinate either
in neutral or deprotonated form,[1] they can be used as
antitumor agents,[2] anticonvulsants[3] and herbicides.[4]
It was found that 2-hydroxy-1-naphthaldehyde
thiosemicarbazone and its derivatives can be used as
selective agents with cholesterol, glucose, and uric and
amino acids in fluorometric determinations, in addition
to Fe(II) and Cu(II) in spectrophotometric techniques.[5]
Also, a large number of thiosemicarbazone compounds
showed medicinal applications owing to their potentially
valuable biological properties. Metal chelates that have
the capacity to join with nucleobases, DNA fragments,
amino acids, peptides, and proteins deserve special atten-
tion owing to the clinical use of transition metal chelates
as antitumor drugs.[4,6] The biological activities of metal
chelates vary from those of either metal ions or ligands
2 | MATERIALS AND METHODS
2.1 | Reagents and instruments
2-hydroxy-1-naphthaldehyde, 4-aminoantipyrine, carbon
disulfide, and methyl iodide were obtained from Sigma
Aldrich, BDH, and Merck, and used as received. Elemen-
tal microanalyses (C, H, and N) were performed by the
Microanalytical Unit (Exeter Analytical, Inc. North
Chelmsford, MA 01863, USA) within Cairo University
(Egypt). Metal elemental analysis and chloride ions were
estimated by the usual methods.[12] Infrared spectra were
recorded on a Nicolet Fourier transform infrared (FT-IR)
spectrophotometer (Thermo Fisher Scientific, North
SCHEME 1 Suggested molecular structures of
complexes 1–7

EL-SAWAF ET AL.
3 of 18
Calorina, USA) in the range 4000–400 cm−1. A Varian
Cary 5000 spectrophotometer (Varian Inc., Palo Alto,
California, USA) was employed to perform the absorp-
tion electronic spectra as polycrystalline solid samples.
as electrolyte and a three-electrode cell containing a
3 mm GC electrode as a working electrode, a platinum
wire as an auxiliary electrode, and silver/silver chloride
(Ag/AgCl) as the reference electrode.
1
The H and 13C NMR spectra of the organic ligand were
recorded in deuterated dimethyl sulfoxide (DMSO-d6) on
a Bruker-400 MHz instrument (Bruker Billerica Massa-
chusetts, USA). The fast atom bombardment (FAB) mass
spectra were measured using a gas chromatograpy-mass
spectrometery (GC-MS, DFS) (Varian Inc., Palo alto,
California, USA). The electron paramagnetic resonance
(EPR) spectra for polycrystalline samples of
Cu(II) chelates were recorded on a Varian E-109c spec-
trometer (Varian Inc., Palo alto, California, USA)
equipped with a field modulation unit at 100 kHz. Mea-
surements were effected in the X-band at room tempera-
ture; the microwave power was around 10 mW. Molar
conductivity measurements were made in dimethyl form-
amide (DMF) solution 10−3 M using a type CD6N
Tacussel conductimeter (Mettler Toledo, Ohio, USA).
Thermal gravity-differetial thermal gravity (TG-DTG)
analysis was conducted using a Shimadzu DAT/TG-50
thermal analyzer (Shimadzu, Tokyo, Japan) with a
heating rate of 10ꢀC/min under a nitrogen atmosphere
with a flow rate of 10 drops per minute (dpm) from room
temperature up to 1000ꢀC using platinum crucibles. Mag-
netic susceptibilities were measured at room temperature
by a modified Gouy method by using a Johnson Matthey
magnetic susceptibility balance (Jhonson Matthey Fabri-
cated Equipment, Wayane Pa. USA). The effective mag-
netic moments were calculated from the equation
μeff = 2.84(XMcorrtT)1/2. Melting points were measured
by using Stuart melting point apparatus (Stuart, Cole-
Parmer, Staffordshire, UK). Cyclic voltammetry was car-
ried out using a potentiostat (PST 006) from Voltalab-
Radiometer Analytical with software model voltaMaster
2.2 | Synthesis of (E)-4-(1,5-dimethyl-
3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-
1-(3-hydroxynaphthalen-2-yl)methylene)
thiosemicarbazide (H2L)
The organic ligand was synthesized as shown in
Scheme 2. 4-(2,3-dimethyl-5-oxo-1-phenyl-2,5-dihydro-
1H-pyrazol-4-yl)thiosemicarbazide was synthesized fol-
lowing the previously reported method.[13] (E)-4-(1,-
5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-1-
(3-hydroxynaphthalen-2-yl)methylene) thiosemicarbazide
(H2L)
ligand
was
prepared
by
reacting
4-(2,3-dimethyl-5-oxo-1-phenyl-2,5-dihydro-1H-pyrazol-4-yl)
thiosemicarbazide (1.65 g, 12.40 mmol) with 2-hydroxy-
1-naphthaldehyde (2.00 g, 12.40 mmol) in 70 ml of absolute
ethanol. The reaction mixture was kept under
reflux at 80ꢀC for 24 hr and monitored by thin-layer
chromatography. On completion, the yellow product was
filtered off, washed several times with cold ethanol, and air
dried under ambient conditions.
2.3 | Synthesis of metal(II) chelates of the
Schiff base ligand
Forty milliliters of an ethanolic solution containing
2.72 mmol of the appropriate metal salt was added
portionwise to a suspension of the thiosemicarbazone
ligand (1.20 g, 2.72 mmol) in ethanol (50 ml) while stir-
ring with gentle warming, and then the reaction mixture
was refluxed at 80ꢀC for 10 hr. The resulting color che-
lates were filtered, washed with anhydrous diethyl ether,
and dried in a vacuum desiccator over anhydrous cal-
cium chloride.
4
(Hach, SelAgua, Marlboug, Wiltshire, UK). All
voltammograms were recorded at 100 mV/s scan rate
for the potential range from 2.0 to −0.5 V in DMF con-
taining 0.1 M tetrabutylammonium perchlorate (TBAP)
SCHEME 2 Preparation of the organic
ligand H2L

4 of 18
EL-SAWAF ET AL.
2.4 | Theoretical calculations
details of the procedure and estimation of the inhibition
zones are described elsewhere.[24]
Geometry optimization and frequency calculations of the
stable conformers of (E)-4-(1,5-dimethyl-3-oxo-2-phenyl-
2,3-dihydro-1H-pyrazol-4-yl)-1-(3-hydroxynaphthalen-
2-yl)methylene) thiosemicarbazide (H2L) were performed
using density functional theory (DFT) at the
B3LYP/LanL2DZ level of theory.[14] The calculated vibra-
tional modes were scaled by a factor of 0.9679.[15] The
maximum electronic absorption bands, vertical electronic
excitations, and oscillator strengths (f > 0 for allowed
transition) of the ligand were calculated using the time
dependent density functional theory (TD-DFT) method
at the same level of theory.[16,17] NMR magnetic isotropic
shielding tensors (σ) were predicted using the gauge-
independent atomic orbital (GIAO) approach[18] using
the above-mentioned hybrid functionals. The isotropic
shielding values were used to calculate the isotropic
chemical shifts δ with respect to tetramethylsilane
(Si(CH3)4). δiso(X) = σTMS(X) – σiso(X), where δiso is the
isotropic chemical shift and σiso the isotropic shielding
constant. The predicted chemical shifts were obtained
using the equation δexp = aδcal + b, where δcal = δiso. The
theoretical calculations in the solvent were obtained
using a polarizable continuum model (PCM) in which
the solute was embedded into a cavity surrounded by a
dielectric continuum solvent described by its dielectric
constant (e.g., εCDCl3 = 4.7113).[19] The PCM has been
reported to correctly model major solvent effects such as
the electrostatic effects of the medium providing no spe-
cific solute–solvent interactions such as hydrogen bond
interactions, dipole–dipole interactions, or induced
dipole–dipole interactions are considered.[20] Theoretical
calculations were performed using the Gaussian09
package.[21]
3 | RESULTS AND DISCUSSION
3.1 | Elementary analyses and structural
studies
The analytical data and physical properties of the ligand
(H2L) and its metal chelates are shown in Table 1. The
results show that the ligand coordinates to the metal ion
in a 1:1 molar ratio. Chelates of the monobasic ligand
(i.e., loss of phenolic OH proton) are formed with the
chlorides of Co(II) and Cu(II), Co(II) acetate,
Cu(II) acetate, and Cu(II) perchlorate. Chelates of the
dibasic ligand (i.e., loss of OH and N(2)H protons) are
formed on chelation with Co(II) nitrate and
Ni(II) chloride. All synthesized solid chelates are colored,
stable to air and moisture, and soluble in DMF and
dimethyl sulfoxide (DMSO). The brown, green, red, and
blue colors are common to chelates involving
thiosmeicarbazone coordination due to the sulfur to
metal charge transfer bands, which dictate their visible
spectra.
The molar conductance of the metal chelates was
measured in DMF solution (10−3 M) and their values
(ΛM) are shown in Table 1. The low molar conductance
values of the chelates confirm their nonelectrolytic
nature.[25–27] Our effort to grow crystals suitable for X-ray
structural determination was unsuccessful.
1
3.2 | H and 13C NMR spectra
One of the most important signs for the formation of H2L
1
was obtained from the H NMR spectrum, where assign-
ments of the signals were established on the chemical
1
2.5 | In vitro antibacterial activity
shifts and intensity patterns of the H NMR spectrum of
the thiosemicarbazone ligand (Supporting Information
Figure S1 and Table 2). The numbering of the ligand is
displayed in its optimized structure obtained at the
B3LYP/LanL2DZ level of theory (Figure 1).
In vitro antibacterial screening of the thiosemicarbazone
(H2L) and its metal chelates was performed at the
Department of Parasitology and Animal Diseases,
National Research Center, Dokki, Egypt. The bacterial
cultures used in this study were Staphylococcus aureus as
representative for Gram-positive bacteria and Escherichia
coli as representative for Gram-negative bacteria using
the agar diffusion method.[22,23] In all measurements, the
concentration of the investigated compounds was
50 mg/ml using dimethyl sulfoxide (DMSO) as solvent
control. Dimethyl sulfoxide (DMSO) was used as solvent
control and standard enrofloxacin antibiotic discs as used
as a positive control in the antibactrial testes. Complete
The downfield shift of the phenolic OH proton which
resonates as singlet signal at δ (ppm) = 11.77 (s, 1H, O–
H), in the ligand spectrum indicates the formation of
strong intramolecular hydrogen bonding through the
N(1) nitrogen atom[28] and confirms that the potential
donors O and N(1) are found in a syn configuration.[29]
Also, the ligand exhibits N(2)H at 10.48 ppm, N(4)H at
9.25 ppm, azomethine proton, N=C(H), at 9.16 ppm, aro-
matic naphthyl and an N-phenyl group of protons as
multiplets in the range 8.54–7.22 ppm, N–C(H3) at

EL-SAWAF ET AL.
5 of 18
3.13 ppm, and C–CH3 at 2.52 ppm. The corresponding
predicted chemical shifts of the ligand were determined
at the B3LYP/LanL2DZ level of theory (Table 2). The
correlation curve between the predicted and experimen-
tal 1H NMR chemical shifts is displayed in Figure 2. Rel-
atively good correlation is obtained between the
predicted and experimental 1H NMR chemical shifts with
a correlation coefficient of 97.61%.
The structure of the ligand was evaluated further
using 13C NMR, where signals for different types of car-
bon atoms were present in the organic ligand spectrum
(Figure S2). The signals observed at 177, 152, and
148 ppm are due to C=S, C=N, and C–O of the
thiosemicarbazone moiety, respectively. The spectrum
also shows a strong signal at 162 ppm that may be
assigned to the C=O of the pyrazolone ring.
3.3 | Mass spectra
The formation and purities of the Schiff base ligand
(H2L) and its metal chelates were confirmed using mass
spectrometry. The mass spectrum of the Schiff base
ligand H2L had a molecular ion peak at m/z = 442.5,
which agrees with the molecular weight (M + 2). The
intensity of the peaks due to fragmentation is shown in
Supporting Information Figure S3. Similarly, the mass
spectra of metal complexes 1–7 are shown in Supporting
Information Figure S4, where the molecular ion peaks
can be seen at m/z = 544, 582, 595, 597.3, 603, 592, and
609, respectively. The results obtained from the elemen-
tal analyses and mass spectra of the synthesized Schiff
base ligand and its metal chelates are consistent with
each other, which confirms the proposed molecular for-
mula (Table 1). In addition, the spectra of the com-
pounds showed some additional peaks owing to
molecular cations of various fragments of the metal
chelates.
3.4 | Infrared spectra
The prominent assignments for the IR spectral bands
for the organic thiosemicarbazone ligand and its metal
chelates are shown in Table 3. In principle, the organic
ligand should display thione–thiol tautomerism, –
N(H)–C(=S)– $ N=C(–SH)–, since it contains a thio-
amide –NH–C=S functional group. The absence of any
ν(S–H) band, which usually appears in the
2600–2800 cm−1 region of the IR spectrum of the ligand
(Supporting Information Figure S5), suggests the
thionic nature of the organic ligand in the solid
state.[30,31]

6 of 18
EL-SAWAF ET AL.
TABLE 2
Calculated and experimental 1H chemical shifts (ppm) of
(E)-4-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)-1-(3-hydroxynaphthalen-2-yl)methylene) thiosemi-carbazide(H2L) in
deuterated dimethyl sulfoxide (DMSO-d6)
Predicted
Experimental
1H NMR
H1
7.73
7.55
7.72
7.88
7.47
7.80
8.58
2.59
3.61
7.58
7.67
7.65
7.79
5.80
12.40
7.90
7.56
7.93
8.35
7.34
7.93
9.17
2.52
3.13
7.87
7.53
7.24
10.44
9.26
11.77
H2
H3
H6
H9
H10
H11
H23
H24
H26, H30
H27, H29
H28
H13 (NH)
H16 (NH)
OH
The spectrum of the free ligand showed a broad
medium band at 3430 cm−1 assignable to intramolecular
hydrogen bonding between phenolic O–H and the
N(1) nitrogen atom.[32,33] The absence of this band in the
spectra of all metal chelates provides strong evidence for
coordination through the deprotonated phenolic oxygen
and this causes ν(C–O) to be blue-shifted by 30–80 cm−1
from its position at 1242 cm−1 in the ligand spec-
trum.[34,35] This is also supported by the presence of a
new band around 550 cm−1 in the spectra of the chelates
assignable to ν(M–O).[36] The spectrum of the ligand
shows a medium band at 3175 cm−1 attributable to ν(N–
H).[37] However, this band disappears in the spectra of
[Co(L)(H2O)].H2O.1.5EtOH and [Ni(L)(H2O)].2EtOH
chelates (Supporting Information Figure S5) due to the
change in the tautomerism and subsequent
deprotonation of the =N–NH– group. Consequently, a
new band due to ν(N=C) is seen at ca. 1580 cm−1
.
Other bands were observed at 3040 and 2930 cm−1, and
these can be attributed to υ(C–H) of the aromatic and ali-
phatic vibrations, respectively.[38] A small decrease in the
energy (about 5–20 cm−1) of the ν(C=N) band is indicative
of coordination via the azomethine nitrogen ion.[39,40] This
is further confirmed by the appearance of a new band in
the range 480–430 cm−1assignable to ν(M–N).[41] The
medium intensity band found at 1150 cm−1 owing to ν(N–
N) of the thiosemicarbazone was shifted to a higher fre-
quency in the spectra of all chelates, indicating the increase
in bond strength, and again endorses coordination via the
azomethine nitrogen. A medium intensity band at
865 cm−1 corresponding to ν(C=S) in the spectrum of the
ligand completely disappeared in the spectra of chelates
FIGURE 1 The optimized structure of
ligand (H2L) obtained at the B3LYP/LanL2DZ
level of theory

EL-SAWAF ET AL.
7 of 18
participation of te thiolate sulfur atom through enolization
and subsequent deprotonation.[42,43]
The spectrum of the acetato chelate, [Cu (HL)
(OAc)].2H2O, showed two bands at 1625 and 1425 cm−1
due to νas(COO−) and νsy(COO−), respectively, while the
corresponding bending modes were seen at 670 and
630 cm−1. The large separation Δν = νas – νsy = 200 cm−1
is indicative for the monodentate geometry of the acetate
ligand bonded to Cu (II).[44] The occurrence of coordi-
nated water in chelates 3 and 4 was confirmed by the
presence of bands at ca. 3460, 1570, 955, and 655 cm−1
due to ν(OH), δ(H2O), ρrock(H2O), and ρwagg(H2O),
respectively.[45] The last two modes were absent from the
spectra of the other chelates with water molecules, indi-
cating hydrated rather than coordinated water.[46]
The spectrum of perchlorate chelate (7) showed split
strong bands (ν3) at 1117 and 1049 cm−1, and a strong
band at 620 cm−1 (ν4) (Supporting Information
Figure S5) in agreement with monodentate coordination
of perchlorate.[47] Thus, it may be established that the
ligand behaves as a monobasic[48] or dibasic[49] tridentate
ONS chelating agent coordinating via azomethine nitro-
gen, phenolic oxygen, and thione/thiol sulfur.
FIGURE 2 Correlation curve between predicted and
experimental 1H NMR chemical shifts of the ligand
[Co(L)(H2O)].H2O.1.5EtOH (3) and [Ni(L)(H2O)].2EtOH
(4) (Supporting Information Figure S5), and new bands at
745 and 755 cm−1 owing to υ(C–S) indicated the
TABLE 3
Important IR spectral bands (cm−1) of the ligand (H2L) and its metal chelates
ν[N(2)
H]
ν(C–O)
ν(C=N) phenolic
ν(M–
ν(C=S) ν(C–S) O)
ν(M–
N)
ν(M–
S)
No. Ligand/chelate
ν(N–N)
960(m)
985(w)
H2L·0.5H2O
3175(m) 1600(s)
3180(w) 1595(s)
1250(s)
865(m)
850(m)
–
–
–
–
–
–
1
2
3
[Co (HL)Cl].H2O
1185(m)
575(w)
490(w)
425(w)
415(m)
418(m)
[Co (HL)(OAc)].2H2O
3100(m) 1597(m) 1190(m)
1020(m) 825(w)
540(m) 460(w)
450(w)
[Co(L)(H2O)].
H2O.1.5EtOH
–
1605(s)
1180(m)
990(s)
–
745(m) 435(w)
4
5
[Ni(L)(H2O)].2EtOH
[Cu (HL)Cl].4H2O
–
1600(w)
1180(m)
1210(m)
995(s)
–
755(s)
560(m) 465(s)
535
(m)
425(w)
3150(m) 1585(s)
1025(s)
815(m)
–
440(m) 422(w)
6
7
[Cu (HL)(OAc)].2H2O
[Cu (HL)(ClO4)].H2O
3120(w)
3110(w)
1590(m) 1195(m)
1605(s) 1190(m)
1020(w)
995(s)
835(w)
840(m)
–
–
548(w)
550(w)
460(w)
463(w)
418(w)
426(w)
TABLE 4
Some of the calculated and predicted vibration modes of the ligand (H2L) and its [Ni(L)(H2O)].2EtOH (4) and [Cu (HL)
Cl].4H2O (5) chelates
H2L
Cal
4
5
Scal
3399
1558
1211
1114
Exp
3175
1600
1250
960
Cal
3582
1610
–
Scal
3467
1558
–
Exp
Cal
3537
1615
–
Scal
3423
1563
0
Exp
3150
1585
1210
1025
νN-H
νC=N
νC-O
νN-N
3512
1610
1251
1151
–
1600
1180
995
–
–
–
0
Note. Cal, Calculated; Exp, Experimental; Scal, Scaled.

8 of 18
EL-SAWAF ET AL.
TABLE 5
Magnetic moment and solid state electronic spectral data (nm) of the ligand and its metal chelates
Compound
π ! π*
220 245
260
n ! π*
280 350
300 343
315 350
305 335
305 355
305 350
310 380
315 355
LMCT
d ! d
μeff(B.M.)
–
H2L
–
–
[Co (HL)Cl]· H2O
[Co (HL)(OAc)]· 2H2O
[Co(L)(H2O)].H2O.1½EtOH
[Ni(L)(H2O)].2EtOH
[Cu (HL)(Cl]·4H2O
[Cu (HL)(OAc)] ·2H2O
[Cu (HL)(ClO4]· H2O
395 440
370 450
390 430
390 445
395 455
395 455
390 450
470 640
485 660
430 615
470 660
465 620
480 645
460 610
4.42
4.26
4.35
2.61
1.75
1.89
1.81
275
280
245
275
275
270
Note. B.M., Bohrmagneton; LMCT, Ligand metal charge transfer.
Some of the calculated and predicted vibration modes
of the ligand and its Ni(II) and Cu(II) chelates are dis-
played in Table 4. The calculated values are obtained at
the B3LYP level of theory and the predicted values are
obtained from the calculated and scaled by a factor of
0.9679. As shown in Table 4, ν(C=N) is relatively well
predicted with deviations of 42, 42, and 22 cm−1 with
respect to the experimental values of the ligand, and its
metal chelates 4 and 5, respectively.
assignable to the phenol (OH) function.[51] The maxi-
mum absorption band at 350 nm is overestimated at the
B3LYP level of theory in gas and PCM solvent with a var-
iation of 26 nm with respect to the experimental value.
This band corresponds to the HOMO ! LUMO elec-
tronic transition (Figure 3).
The absorption bands at 245 nm are attributed to elec-
tronic transitions HOMO-7
!
LUMO (37%) and
HOMO ! LUMO+4. This band is overestimated by
14 nm with respect to the experimental value. These tran-
sitions are also found in the spectra of the metal chelates,
but they are shifted towards lower and higher frequen-
cies, confirming coordination of the ligand by the metal
ions.[52] Based on previous studies of the metal chelates
of thiosemicarbazones,[53,54] we assigned a π ! π* band
at ca. 270 nm and an n ! π* band at ca. 300 nm
(Table 5). The shift of the n ! π* bands to longer wave-
length region in metal chelates is strong evidence for the
coordination of azomethine nitrogen and thioamide sul-
fur with metal ions.[55] To emphasize the chelation effect
on the absorption bands observed in the ligand spectrum,
the predicted spectra of the metal chelates 2, 4, and 5 were
obtained at the B3LYP level of theory (Figure 4). As can
be seen in Figure 4, the chelation of the ligand with metal
ions induces a bathochromic shift of π ! π*. The
observed bands at 249 and 276 nm were shifted to
345 and 396 nm, respectively. From the predicted spectra
of the ligand and its Ni chelate 4, it is clear that the
3.5 | Magnetic moments and electronic
spectra
The magnetic moments values of the metal chelates at
room temperature in accordance with the electronic spec-
tral data of the free ligand (H2L) and its metal chelates
are listed in Table 5. The spectra of the organic ligand
(H2L) and its metal chelates were measured in Nujol
mull over the range of 200–1000 nm.
The spectrum of the free thiosemicarbazone
(Supporting Information Figure S6) shows four absorp-
tion bands at 220, 245, 280, and 350 nm. The first two
bands are assigned to the intraligand π ! π* transitions
while the third band at 280 nm is assigned to n ! π*
charge transfer transition associated with the azomethine
and thioamide functional groups overlapping in the same
envelope.[50] The n ! π* transitions at 350 nm are
FIGURE 3 HOMO (left) and LUMO (right) frontier orbitals of the ligand H2L

EL-SAWAF ET AL.
9 of 18
FIGURE 4 Predicted spectra of the
ligand H2L and its complexes 2, 4, and 5
chelation leads to a bathochromic redshift of n ! π*.
Similarly, the Cu chelate 5 with the ligand leads to a bat-
hochromic shift of π ! π*. The HOMO and LUMO
orbitals are shown in Figure 5. The HOMO and LUMO
orbitals of the 2, 4, and 5 metal chelates are displayed in
Figure 4. As can be seen from Figure 5, the
HOMO ! LUMO transition induces a transfer of charge
from one side to another in these chelates. For instance,
in 4, the HOMO ! LUMO transition induces a transfer
of charge from the naphthalene moiety to Cu and its vici-
nal moieties (Figure 5).
The metal to ligand charge transfer bands are found
at 350–400 and 400–450 nm.[56] The lower wavelength
band is assigned to S ! M(II) transitions, which are quite
common in the electronic spectra of metal chelates of
thiosemicarbazones,[57] and the higher wavelength band
is associated with N ! M(II) transitions. It has been
shown that the phenolate compounds show a broadening
line with a tail running across the visible part of the spec-
trum resulting from phenolate to metal charge transfer
that overlaps on the high wavelength side of
S ! M(II) transitions. For the metal chelates under
study, the O ! M(II) transitions are found in the range
290–350 nm.[58]
The electronic spectra of the dark green and blue
Co(II) compounds show two bands at ca. 666 and 465 nm
4
which are assigned to A2g
!
4T1g(F) (υ2) and
4A2g
!
4T1g(P) (υ3) transitions, respectively, revealing
typical tetrahedral coordination around the Co(II) ion.[59]
The values of ligand field parameters B (interelectronic
repulsion of the
d electrons in chelate), β (the
nephelauxetic effect), and 10Dq are calculated as stated

10 of 18
EL-SAWAF ET AL.
FIGURE 5 HOMO (left) and LUMO (right)
frontier orbitals of metal complexes 2, 4, and 5
by the equations reported for the tetrahedral compounds
1, 2, and 3, and their values are found to be in the ranges
reported for other tetrahedral Co(II) chelates with sulfur
donor atoms.[60,61]
The μeff values 4.42, 4.26, and 4.35 B.M. for
Co(II) chelates 1, 2, and 3, respectively, correspond to
three unpaired electrons in a tetrahedral environment.[64]
The electronic spectrum of the Ni(II) chelate 4 shows two
d–d bands at 470 and 660 nm, the position of the bands
indicative of tetrahedral stereochemistry, assigned to
3T1(F) ! 3T2 (υ2) and 3T1(F) ! 3T2 (υ3) transitions,
respectively. More intense bands in the d–d electronic
spectrum and high magnetic moment value, μeff = 2.61 B.
μeff = 3:87ð1 – 4λ=10DqÞ,λ = – 178 cm– 1
Â
Ã
2
B = 4ðν3 – 15DqÞ – 10Dq2 =½60ðν3 – 15DqÞ – 180Dqꢁ
M., support
a
tetrahedral configuration for the
ν1 = 10Dq
Ni(II) chelate.[65] Cu(II) chelates in the present study
show μeff values of 1.75, 1.89, and 1.81 B.M. for chelates
5, 6, and 7, respectively, corresponding to one unpaired
electron. The absorption spectra of Cu(II) chelates are
consistent with the earlier report of square planer che-
lates[46] and show a low intense d–d band in the range
Β = BðchelateÞ=Bðfree ionÞ
where B(free ion) for Co(II) is 963 cm−1
.
The 10Dq values obtained for metal chelates 1–3 via
this method are −5170, −5140, and −4980 cm−1, respec-
tively. These results indicate that the interelectronic
repulsion of the d electrons in metal chelates is less than
in the free ion. The B values for the compounds are about
70% of the free ion value. The reduction of B by chelate
formation is affected by the delocalization of the d elec-
tron cloud on the organic thiosemicarbazone, which is
sequentially caused by the formation of the covalent
bond.[62] According to Jorgenson, the organic ligand
which coordinates via a sulfur donor atom shows a low
value of B and causes a pronounced nephelauxetic
effect.[63] The lower the value of the nephelauxetic effect,
β, the greater the extent of covalence. For Co(II) chelates,
β values are lower than 1, which provides the informa-
tion on the covalent nature of the metal–ligand bond.
2
2 2
600–650 nm corresponding to 2B1g ! Eg(dx –y ! (dxz,yz),
2
2
2
2
2
2
2
2
B1g ! B2g (dx –y ! dxy), and 2B1g ! A1g (dx –y ! dz ).
Since these lower orbitals are often close in energy, it is
difficult to resolve the individual transitions from the
lower to upper d level and a single absorption band was
observed. 66]
[
3.6 | Thermal studies
The thermal pyrolysis behavior of the free
thiosemcarbazone ligand and its metal chelates were
examined by thermogravimetric analysis (TG) and differ-
ential thermogravimetry (DTG) in the temperature range
of room temperature up to 1000 ꢀC, under an inert

EL-SAWAF ET AL.
11 of 18
atmosphere. The percentage of weight losses found
(calcd), temperature range, and interpretation of each
pyrolytic step are presented in Table 6 and thermograms
are shown in Supporting Information Figure S7. Because
of the explosive nature of perchlorate salts, the thermal
properties of chelate 7 were not determined.
The thermoanalytical data in Table 6 show that the
organic ligand (H2L) exhibited a weight loss of 2.34%
within the range 25–215ꢀC equivalent to a loss of 0.5 mol
of water. The ligand undergoes thermal pyrolysis via two
degradation steps. The first occurs at 215–390ꢀC and is
characterized by a weight loss of 32.03%. This weight is
consistent with the elimination of the C10H7O portion via
cleavage of the longest C(7)–C(11) bond in the
thiocarbohydrazide backbone. The second weight loss of
59.91% appeared in the range 395–1000ꢀC compatible
with the liberation of the C11H14N5OS fragment. This
step is associated with a sharp and broad DTG peak at
Tmax = 485 and 610ꢀC, respectively. The weight percent-
age of 5.72 (calcd 5.45%) of the final residue indicates that
the product of the thermal degradation processes of the
thiosemcarbazone ligand is 2 moles of carbon.
The thermal pyrolysis data (Table 6) of the tetra-
coordinated chelates [Co (HL)Cl].H2O (1), [Co (HL)
(OAc)].2H2O (2), [Cu (HL)(Cl].4H2O (5), and [Cu (HL)
(OAc)].2H2O (6) are similar, indicating that they are
TABLE 6
Thermal decomposition of the ligand (H2L) and its metal chelates
Mass loss %
TG range
(ꢀC)
DTG peak
(ꢀC)
No. Chemical formula
Found Calcd Assignment
H2L.0.5H2O C23H22N5O2.5
S
25–215
–
285
2.34
32.03
59.91
5.72
2.05
32.50
60.00
5.45
Loss of 0.5H2O
215–390
390–1000
At 1000
25–100
Loss of C10H7O
Loss of C11H14N5OS
2 Cc
485, 610
1
2
3
[Co (HL)Cl].H2O C23H22N5O3SCoCl
[Co (HL)(OAc)].2H2O C25H27N5O6SCo
–
198
3.32
3.54
Loss of H2O moleculesa
100–220
195–1000
At 1000
25–100
6.53
6.22
Loss of 0.5Cl2
266, 371
61.96
28.19
6.16
61.59
28.66
6.48
Loss of C16.5H20N5OS
CoO + 6.5Cc
Loss of 2H2O moleculesa
Loss of CH3COO
Loss of C18H20N5OS
CoO + 5Cc
39,80
202
100–220
220–1000
At 1000
25–180
10.11
60.64
23.09
17.71
10.26
60.21
23.05
17.65
269, 375
[Co(L)(H2O)].H2O.1.5EtOH
C26H32N5O5.5SCo
145
256
Loss of 2H2O + 1.5
EtOHa
180–1000
At 1000
25–200
69.67
12.62
18.41
69.45
12.14
11.98
5.89
69.85
12.50
18.22
69.10
12.68
11.55
5.13
Loss of C23H19N5OS
CoOc
Loss of H2O + 2 EtOHa
Loss of C23H19N5OS
Nic
4
5
[Ni(L)(H2O)].2EtOH C27H33N5O7Sni
[Cu (HL)Cl].4H2O C23H22N5O7SCuCl
143
270
–
200–1000
At 1000
25–150
98
Loss of 4H2O moleculesa
150–250
250–1000
At 1000
25–110
177
274
Loss of 0.5Cl2
52.93
29.20
6.12
53.48
28.84
5.77
Loss of (C15H20N5O2S)
CuO + 8Cc
Loss of 2H2O moleculesa
Loss of CH3COOb
Loss of C15.25H20N5OS
CuO + 7.75Cc
6
[Cu (HL)(OAc)].2H2O C25H27N5O6SCu
58
110–250
250–1000
At 1000
180(b)
10.02
54.56
29.30
10.34
54.52
29.37
Note. DTG, Differential thermal gravity; TG, thermal gravity.
aDehydration.
bDesolvation.
cFinal residue.

12 of 18
EL-SAWAF ET AL.
TABLE 7
Cu(II) chelates 5–7
Electron spin resonance spectral parameters of the
isothermal. The theromgrams showed weight losses
within 25–100ꢀC due to the dehydration process. The
compounds [Co (HL)Cl].H2O (1) and [Co (HL)
(OAc)].2H2O (2) started decomposition at 100ꢀC, while
the chelates [Cu (HL)(Cl].4H2O (5) and [Cu (HL)
(OAc)].2H2O (6) begin to decompose at 110 and 150ꢀC,
respectively. The decomposition process was achieved
mainly through two processes. The first step was TG
weight losses corresponding to the removal of coordi-
nated anions in the temperature range 100–220ꢀC as
shown in Table 6. The thermograms for these compounds
Parameter
Chelate 5
2.214
Chelate 6
Chelate 7
g///g3
2.219
2.225
g2
2.162
g⊥/g1
2.041
2.048
2.105
4.742
2.051
2.109
4.551
a
gav
go = 2.139
2.123
G
A// × 10−4 cm–1
A⊥ × 10−4 cm–1
174
revealed
a varying weight loss within the range
Δ2
Δ3
k2==
k⊥2
K
α2
β2
γ2
f
16129
21505
0.515
0.531
0.73
15503
20833
0.507
0.575
0.55
16393
21739
0.551
0.639
0.61
~220–1000ꢀC corresponding to organic ligand pyrolysis,
leading to M(II) oxides with contamination with variable
amounts of carbon atoms as a final product of the ther-
mal decomposition process.
The thermoanalytical data and thermograms of the
tetrahedral chelates [Co(L)(H2O)].H2O.1.5EtOH (3) and
[Ni(L)(H2O)].2EtOH (4) decompose in a similar way
(Supporting Information Figure S7). The thermal degra-
dation processes occurred via two distinct pyrolytic
steps. The first steps are accompanied by weight losses
of 17.71 and 18.41% within the range 25–200ꢀC for
3 and 4, respectively, consistent with the elimination of
ethanol and coordinated water molecules in one
step.[67] The second steps are characterized by sharp
DTG peaks at Tmax. = 256 and 270ꢀC and weight losses
of 69.67% and 69.45% for 3 and 4, respectively, due to
complete ligand pyrolysis leaving CoO and NiO as final
residue.
0.57
0.90
0.93
127
agav = (2 g⊥ + g//)/3, K = (2K⊥ + K//)/3.
of g3 > g2 > g1 suggests the anisotropy of the chelate. The
trend g// > g⊥ > 2.0023 shown for all chelates (for g ten-
sors
with
orthorhombic
splitting
we
take
g⊥ = (g1 + g2)/2[68]) shows that the unpaired electron
2
2
2
most likely resides in the dx –y orbital, implying B1g as
the ground state, which is characteristic of square planar
geometry.[69,70] For chelates with dx –y ground state,
strong interaction along the z axis is complemented by an
increase in the value of g//. Strong axial bonding causes
an increase in the length of the bond in the x–y plane
which leads to a decrease in both in-plane covalency and
2
2
3.7 | Electron Spin Resonance (ESR)
spectra of Cu(II) chelates
The X-band electron spin resonance spectra of the poly-
crystalline thiosemicarbazone Cu(II) chelates 5–7 were
recorded at room temperature and the values of the spin
Hamiltonian parameters of the chelates were calculated
and are presented in Table 7.
The shape of the electron spin resonance spectra of
the chelates is consistent with square planar geometry
around each Cu(II) center in the chelates, which agrees
with the electronic spectral data.
The electron spin resonance spectra of chelates 5–7
show two different types of geometrical species. Com-
pounds 6 and 7 show a typical axial spectrum with well-
defined parallel and perpendicular spin (g// and g⊥) being
assignable. Spectrum of the chelate 5 shows three reso-
nance signals and the effective g values were found to be
g1 = 2.041, g2 = 2.162, and g3 = 2.214. The values of g1
and g2 for compound 5 are very close, indicating that the
rhombic distortion is very small and the decreasing order
the energy of the dx –y2 transition. Both these effects tend
2
to increase the value of g//. Kivelson and Neiman have
shown that g// is the most sensitive function for rep-
resenting covalency.[71] For the ionic environment, the
g// value is normally 2.3 or higher and less than 2.3 for
the covalent character to the metal-ligand bonds in che-
lates. It is obvious from Table 7 that the g// obtained is
less than 2.3, indicating the covalent character of the
metal–ligand bond in the existing chelates. Massacesi
et al. reported that g// is 2.3–2.4 for Cu–O bonds, 2.2–2.3
for Cu–N bonds and mixed Cu–N and Cu–O systems,
and 2.1–2.2 for Cu–S bonds.[72] For the current chelates
g// is 2.214–2.225, confirming the presence of Cu–O, Cu–
N, and Cu–S bonds in these compounds. The parameter
R (R = (g2 – g1)/(g3-g2) for rhombic systems) calculated
2
2
for the compound 5 is 0.70, i.e., R < 1, indicating a dx –y
ground state for the Cu(II) ion.[73]

EL-SAWAF ET AL.
13 of 18
TABLE 8
Electrochemical data for the ligand and its metal
The axial geometric parameter G, which is a measure
chelates
of the exchange interaction between the Cu(II) centers in
the polycrystalline solids, is calculated using
G = (g// – 2.0023)/(g⊥ – 2.0023) for axial spectra and
G = (g3 – 2.0023)/(g⊥ – 2.0023), where g⊥ = (g1 + g2)/2,
for rhombic spectra. According to Hathaway and
Billing,[68] if G > 4, the exchange interaction between
Cu(II) centers in the solid state is negligible, but if G < 4
there is considerable exchange interaction in the solid
ΔEp
Compound
Epa
Epc
–
E1/2
–
(mV)
H2L
1.33
1.55
0.66
0.75
–
–
–
–
–
–
[Co(HL)Cl].H2O
–
–
–
–
state chelates. The calculated
G values for the
[Co(HL)(OAc)].2H2O
1.32 1.18 1.250 140
1.54 1.41 1.475 130
Cu(II) chelates are given in Table 7 and suggest that there
is no exchange interaction in solid chelates except for
chloro compound 5.
[Co(L)(H2O)].
H2O.1½EtOH
0.85 0.78 0.815
70
–
1.19
1.38
–
–
–
–
The spectrum of chelate 5 (Supporting Information
Figure S8), display a hyperfine anisotropy electron spin
resonance spectrum with the defined hyperfine structure
on the lower field region owing to the interaction
between the unpaired electron of the Cu(II) ion with its
nucleus, where the values of g and A are used to deter-
mine the empirical factor f = g/A, which is used as an
index for tetragonal distortion.[74] A value of f in the
range 105–135 cm indicates square planar geometry
whereas larger values indicate tetragonally distorted che-
lates.[75] The f value of the investigated chelate (f = 127)
indicates square planar geometry.
–
[Ni(L)(H2O)].2EtOH
1.01 0.84 0.925 170
1.31
1.52
1.05
1.35
–
–
–
–
–
–
–
–
–
–
–
–
[Cu(HL)(Cl].4H2O
[Cu(HL)(OAc)].2H2O
[Cu(HL)(ClO4].H2O
0.91 0.65 0.780 260
1.28
0.62
–
–
–-
Note. ΔEp = (Epa – Epc) where Epa and Epc are anodic and cathodic
potentials, respectively; E1/2 = 0.5(Epa + Epc); scan rate: 100 mV/s.
The bonding factors α2, β2, and γ2, which can be used
as a measure of the covalence of the in-plane σ bonds,
and in-plane and out-of-plane π bonds, respectively, were
calculated. α2 was calculated by the following
covalency of the in-plane π bonding and out-of-plane π
bonding, respectively. Hathaway[68] states that, for pure σ
bonding K ffi K⊥ ffi 0.77 and for in-plane π bonding
K < K⊥, while for out-of-plane π bonding K > K⊥K⊥. For
chelate 5, K < K⊥, indicating considerable in-plane π
bonding.[77] Moreover, lower values of K than unity
equation[71,76]
:
α2 = A=0:036 + ðg−2:0023Þ + 3=7ðg⊥ −2:0023Þ + 0:04:
If the value of α2 = 0.5, this indicates a complete cova-
lent character, while a value of α2 = 1 suggests complete
ionic bonding. The calculated value of α2 for the [Cu(HL)
(ClO4)]chelate is 0.54, representing a portion of unpaired
electron density located on the Cu(II) ion.
The orbital reduction factors K and K⊥ (K2 = α2β2 and
K2 = α2γ2), which are the parallel and perpendicular
⊥
components of the orbital reduction factor and give sig-
nificant information about the nature of the bonding in
Cu(II) chelates, can be calculated using the following
expressions[67]
:
K2 = ðg−2:0023ÞΔE2=8λ0
K2⊥ = ðg⊥ – 2:0023ÞΔE3=2λ0
where λo is the spin-orbit coupling constant of the free
2
2
Cu(II) ion (λ0 = −828 cm−1) and ΔE2 = B1g ! A1g and
FIGURE 6 Cyclic voltammogram of the ligand H2L at scan
rate 100 mV/s
2
2
ΔE3 = B1g ! Eg may be used to measure the degree of

14 of 18
EL-SAWAF ET AL.
(0.55–0.75) are revealing of their covalent environ-
ment.[78] In addition, the values of β2 and γ2 (0.76 and
1.27, respectively) for metal chelate 5 reflect the presence
of in-plane π bonding and ionic character of out-of-plane
π bonding.[79]
bound to the ligand, which indicates that the metal cen-
ter has an equal attitude towards donor atoms.[83]
Ni(II) chelate 4 shows two irreversible oxidation
peaks at Epa = 1.31 V and Epa = 1.52 V corresponding to
the ligand moiety (Figure 7). The chelate also exhibits
one quasi-reversible redox wave, indicating that Ni(II)
chelate undergoes Ni(II)/Ni(III) oxidation at Epa = 1.01 V
and corresponding reduction at Epc = 0.84 V.[84]
3.8 | Electrochemical studies
Extensive studies were reported on four coordinate
Cu(II) chelates.[79,85] Chelate 6 with a coordinated acetate
ligand showed one quasi-reversible Cu (II)–Cu(III) redox
couple at E1/2 = 0.78 V with a peak-to-peak separation
(ΔEp) of 260 mV. Chelates 5 and 7 with coordinated
chloro and perchlorate ligands only showed oxidation
responses at Epa = 1.05 V and Epa = 0.62 V, respectively,
which were attributed to irreversible electrochemical
behavior (Figure 7).
Generally, the nature of the central metal atom
weakly affects the values of the oxidation potentials of
these chelates. However, in comparison with the free
ligand HL, we note a sensitivity of the potential values of
the oxidation of thiosemicarbazone groups to the coordi-
nation of the azomethine, thione/thiol, and phenolate
fragments to the transition metal ion.
The electrochemical properties of the synthesized ligand
and its chelates were studied by cyclic voltammetry mea-
surement in the range of 2.0 to −0.5 V. The voltammetric
data are presented in Table 8 and the values shown are in
agreement with the reported values for redox processes of
similar ONS chelates.[79] The observed cyclic
voltammograms are characterized by either irreversibility
or quasi-reversibility as shown by wide separation of
anodic and cathodic peaks, as well as the absence or mar-
ked reduction in the peak height of the reverse scan.
The voltammogram of the ligand (Figure 6) shows
two irreversible anodic peaks at Epa = 1.33 V and
Epa = 1.55 V, which were assigned to oxidations at the
azomethine and thione groups of the thiosemicarbazone
moiety.[80,81] The shift of redox peaks to a different poten-
tial in the spectra of the chelates indicates the coordina-
tion of the ligand with metal ions where most chelates
under study show one quasi-reversible oxidation process
as manifest by the peak-to-peak separation,
ΔEp > 100 mV.[82]
3.9 | Biological activity
The in vitro antibacterial activities of the
thiosemicarbazone ligand and its metal chelates were
assessed against the human pathogenic bacteria Gram-
positiveS. aureus and Gram-negativeE. coli using the disc
diffusion method. The radial growth of the colony was
recorded on completion of the incubation, and the mean
diameter for each compound at a concentration of
5 mg/ml was recorded. Dimethyl sulfoxide (DMSO) was
used as a solvent and the mean inhibition zone was mea-
sured and compared with the standard antibiotic,
enrofloxacin (5 mg/ml). The images of the zone of inhibi-
tion were shown in Diagram 1 and the results are pres-
ented in Table 9, taking into consideration the stability of
the compounds over a range of pH (5–11) to fit their use-
fulness in antibacterial studies.
For the Co(II) chelate 3, there are three oxidation
peaks on the voltammogram (Figure 7). The first peak at
Epa = 0.85 V, which is not shown by the corresponding
ligand, is quasi-reversible (E1/2 = 0.815 V) and specifies
oxidation of Co(II) to Co(III). The positive potential spec-
ifies that the metal in the lower oxidation state is strongly
It can be seen that all the chelates are more effective
bactericides than the ligand. The detected activities of the
metal chelates are analogous with those conveyed for
chelates prepared using thiosemicarbazones as
ligands.[86] This improvement of the activity can be eluci-
dated based on chelation theory.[87–91] Chelation
decreases the polarity of the metal ion significantly,
mostly owing to the partial sharing of its positive charge
with donor groups and possible π-electron delocalization
on the entire chelate ring. Chelation is not the single
FIGURE 7 Cyclic voltammograms of complexes 3, 4, and 6

EL-SAWAF ET AL.
15 of 18
DIAGRAM 1 The zone effect of the ligand
H2L and its complexes
and pharmacokinetic factors also have a significant role
in determining the effectiveness of an antimicrobial
agent. A way from this, the mode of action of these com-
pounds may also invoke hydrogen bonding between the
>C=N–N=CS– group and the active centers of biomole-
cules and therefore interfere with normal cell processes.
A heterocyclic ligand with many donor atoms has a
greater chance of interaction either with nucleoside bases
or with the essential metal ions present in the biosystem,
so it can be used as bactericides since they tend to inter-
act with some enzymatic functional groups to achieve
higher coordination numbers.[92,93] Furthermore,
enhancement in antibacterial activity on coordination
with a metal ionmay due to the existence of donor atoms
in the ligand or the formed unsaturated coordinately spe-
cies which inhibit enzyme production.[94] All metal che-
lates under study exhibited better activity than the
thiosemicarbazone ligand (H2L) against the tested strains
of bacteria. Furthermore, the free ligand and its metal
chelates were more active against Gram-positive bacteria
(S. aureus) compared to Gram-negative bacteria
(E. coli).[95] In Gram-positive bacteria, chelates 4, 5, and
6 were highly active with comparable zones of inhibition
of 14.3, 14.8, and 14.6 mm, which is almost two times bet-
ter than the ligand. However, in Gram-negative bacteria
metal chelate 5 was the most active, screening a compa-
rable zone of inhibition of 16.1 mm, displaying better effi-
ciency than the ligand and the other chelates. It is worth
TABLE 9
thiosemicarbazone ligand and its metal chelates
In vitro antibacterial activity of tested
Microorganism inhibition zone
diameter (mm)
Gram-positive
bacteria
Gram-negative
bacteria
Staphylococcus
aureus
Compounds
Escherichia coli
HL
1
8.1
10.0
9.8
7.6
11.0
8.7
2
3
8.6
8.3
4
14.3
14.8
14.6
11.6
–
13.7
16.1
13.9
10.1
–
5
6
7
Dimethyl sulfoxide
(DMSO)
Enrofloxacin
26.3
–
principle for antibacterial activity, where some vital fac-
tors such as the nature of the metal ion, the nature of the
organic ligand, coordinating sites, the geometry of the
chelate, concentration, hydrophilicity, lipophilicity, and
occurrence of co-ligands have a significant effect. Steric

16 of 18
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> Ni(II) >
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4 | CONCLUSION
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In this study, Co(II), Ni(II), and Cu(II) chelates with a tri-
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characterized by various analytical and spectroscopic
techniques. Based on physicochemical studies the
Co(II) and Ni(II) chelates adopted tetrahedral geometry
while the Cu(II) chelates demonstrated square planar
geometry as shown in Scheme 1. Furthermore, antimicro-
bial studies showed that the metal chelates have better
inhibitory activity against the tested S. aureus and E. coli
bacteria compared to the unchelated Schiff base
thiosemicarbazone. The inhibitory effect can be corre-
lated with the decrease in polarity of the metal ion, due
to chelation, this could reinforce the lipoliphilicity of the
metal chelates and thus favors its penetration through
the lipid layer of the bacterial membranes. The results
obtained are very interesting and promising, particularly
when compared with the activity displayed by compara-
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ACKNOWLEDGEMENTS
The authors thank Dr. El Hassane Anouar for helping in
the computational part and for useful discussions.
ORCID
Ayman K. El-Sawaf https://orcid.org/0000-0002-0511-
4061
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SUPPORTING INFORMATION
Additional supporting information may be found online
in the Supporting Information section at the end of this
article.
How to cite this article: El-Sawaf AK, Abdel-
Monem YK, Azzam MA. Synthesis, spectroscopic,
electrochemical characterization, density
functional theory (DFT), time dependent density
functional theory (TD-DFT), and antibacterial
studies of some Co(II), Ni(II), and Cu(II) chelates
of (E)-4-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-
1H-pyrazol-4-yl)-1-(3-hydroxynaphthalen-2-yl)
methylene) thiosemicarbazide Schiff base ligand.
Appl Organomet Chem. 2020;e5729. https://doi.
org/10.1002/aoc.5729