Structural characterization, thermal investigation and biological activity of metal complexes containing Schiff Base ligand (Z)-3-(1-((4,6-dimethyl-1H-pyrazolo[3,4-b] pyridin-3-yl)imino)ethyl)-4-hydroxy-6-methyl-2H-pyran-2-one

Article abstract of DOI:10.1002/aoc.4847

A new series of metal complexes containing Co(II), Pd(II), Fe(III) chloride and Cu(II) salts (chloride, bromide, sulphate and perchlorate) have been prepared with Schiff base ligand (HL). The synthesized compounds were elucidated using elemental analyses, spectral techniques, molar conductance, magnetic measurements and thermogravimetric studies. The analytical data established (1?M:1?L) stoichiometry for complexes (1), (2), (4), (6) and (7) as well as (1?M:2?L) and (2?M:3?L) stoichiometry for complexes (5) and (3), respectively. As a result, the ligand HL coordinates in complexes (1), (2), (4), (6) as a monobasic tridentate ONN moiety via the oxygen atom of the deprotonated phenolic OH, the nitrogen atoms of the azomethine and the imine group in pyrazolopyridine ring. While, it behaves as a neutral bidentate in complexes (3,7), chelates via oxygen and nitrogen atoms of enolic OH and azomethine groups. Also, in complex (5) Cu²⁺ ion binds via NO sits of two ligand molecules in its monobasic and neutral forms. The magnetic moment and electronic spectral data proposed octahedral structure for complexes (2, 3 and 7) as well as triagonal bipyramidal and square pyramidal geometry for complexes (1 and 4), while, chelates (5) and (6) possess square planar geometry. TG/DTG studies confirmed the chemical formula for these complexes and established the thermal decomposition processes ended with the formation of metal or metal oxides contaminated with carbon residue. An axial electron spin resonance spectra were suggested for Cu(II) complexes pointing to ²B_1g as a ground state with hyperfine structure for complex (4). In vitro antibacterial and antioxidant activities were performed for HL ligand and its metal complexes. The biological studies indicate that complex (3) has better antibacterial activity compared to the ligand and the other complexes.

Full text of DOI:10.1002/aoc.4847

Received: 2 November 2018
DOI: 10.1002/aoc.4847
Revised: 3 January 2019
Accepted: 14 January 2019
F U L L P A P E R
Structural characterization, thermal investigation and
biological activity of metal complexes containing Schiff Base
ligand (Z)‐3‐(1‐((4,6‐dimethyl‐1H‐pyrazolo[3,4‐b] pyridin‐3‐
yl)imino)ethyl)‐4‐hydroxy‐6‐methyl‐2H‐pyran‐2‐one
Sanaa M. Emam
| Saeyda A. Abouel‐Enein
| Eman M. Abdel‐Satar
Chemistry Department, Faculty of
Science, Menoufia University, Shebin El‐
Kom, Egypt
A new series of metal complexes containing Co(II), Pd(II), Fe(III) chloride
and Cu(II) salts (chloride, bromide, sulphate and perchlorate) have been pre-
pared with Schiff base ligand (HL). The synthesized compounds were eluci-
dated using elemental analyses, spectral techniques, molar conductance,
magnetic measurements and thermogravimetric studies. The analytical data
established (1 M:1 L) stoichiometry for complexes (1), (2), (4), (6) and (7)
as well as (1 M:2 L) and (2 M:3 L) stoichiometry for complexes (5) and
(3), respectively. As a result, the ligand HL coordinates in complexes (1),
(2), (4), (6) as a monobasic tridentate ONN moiety via the oxygen atom of
the deprotonated phenolic OH, the nitrogen atoms of the azomethine and
the imine group in pyrazolopyridine ring. While, it behaves as a neutral
bidentate in complexes (3,7), chelates via oxygen and nitrogen atoms of eno-
lic OH and azomethine groups. Also, in complex (5) Cu2+ ion binds via NO
sits of two ligand molecules in its monobasic and neutral forms. The mag-
netic moment and electronic spectral data proposed octahedral structure for
complexes (2, 3 and 7) as well as triagonal bipyramidal and square pyramidal
geometry for complexes (1 and 4), while, chelates (5) and (6) possess
square planar geometry. TG/DTG studies confirmed the chemical formula
for these complexes and established the thermal decomposition processes
ended with the formation of metal or metal oxides contaminated with carbon
residue. An axial electron spin resonance spectra were suggested for Cu(II)
Correspondence
Eman M. Abdel‐Satar, Chemistry
Department, Faculty of Science, Menoufia
University, Shebin El‐Kom, Egypt.
Email: emanabdelsatar39@gmail.com
2
complexes pointing to B1g as a ground state with hyperfine structure for
complex (4). In vitro antibacterial and antioxidant activities were performed
for HL ligand and its metal complexes. The biological studies indicate that
complex (3) has better antibacterial activity compared to the ligand and the
other complexes.
KEYWORDS
biological studies, metal complexes, Schiff base, spectral techniques, thermal studies
Appl Organometal Chem. 2019;e4847.
wileyonlinelibrary.com/journal/aoc
© 2019 John Wiley & Sons, Ltd.
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https://doi.org/10.1002/aoc.4847

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EMAM ET AL.
(GSK‐3), corticotrophin‐releasing factor (CRF) antago-
nist, antidiabetic, antiarrhythmic. In addition, they are
utilized as anti‐HIV‐1, potential glucocorticoid receptor
ligand for positron emission tomography (PET).[39–41]
1 | INTRODUCTION
Schiff base ligands have been employed intensively in
analytical chemistry because of their excellent ability for
complex formation with different metal ions. Also, this
type of ligands is used as intermediates for the synthesis
of some bioactive compounds. In addition, the ligands
and their complexes depict antibacterial, antifungal,
anticancer activities along with clinical and industrial
uses.[1–11] Schiff bases are well used as antituberculosis,
DNA binding and cleaving agents due to the existence
of chelating azomethine group. On the other hand, the
transition metal complexes are utlized for the synthesis
of cancer combating nonradioactive tools for chemother-
apy and diagnosis along with their uses in supramolecu-
lar photochemistry and in medicine.[12,13]
Dehydroacetic acid (DHA), 3‐acetyl‐6‐methyl‐2H‐
pyran‐2,4(3H)‐dione is one of the oxygen heterocyclic
chelating agents and it is used for the synthesis of wide
variety of heterocyclic ring systems.[14] DHA compounds
have represented a permanent position in coordination
chemistry that explores pharmacological and therapeutic
activities.[15–17] Schiff bases derived from condensation
of DHA with different amines, amino alcohols and amino
phenols have various uses including analytical, clinical,
biological and medicinal fields.[18–20] A number of metal
complexes containing DHA and its Schiff bases have been
reported to exhibit antitumor, antimicrobial, analytical,
medicinal and supramolecular chemistry.[21–26] The bio-
logical activity associated with these chelates is ascribed
to the presence of azomethine group (C=N).[27]
2 | EXPERIMENTAL
2.1 | Material and physical measurements
All chemicals used were chemically pure grade and were
utilized without further purification. CoCl2.6H2O,
CuCl2.2H2O, CuBr2, CuSO4.5H2O, Cu (ClO4)2.6H2O,
PdCl2 and FeCl3.6H2O salts were used as received.
Elemental analyses (C, H, N) were performed on
Perkin Elmer‐2400 elemental analyzer at Main Defense
Chemical Laboratory. All the metal ions of the prepared
complexes were determined with Inductive Coupled
Plasma Mass Spectrometry instrument (ICP), Perkin
Elmer Model: Optima 7000 DV in solution. The solution
of the investigated metal ions were prepared by decompo-
sition (0.08 g) of complexes with concentrated HNO3
mixed with concentrated HCl (v/v: 1:3), then followed
by dilution with distilled water. While, the % of halide
content (chloride or bromide) in the complexes was indi-
rectly estimated by Mohr's method[42] after decomposition
of the samples in concentrated nitric acid by boiling,
followed by addition of distilled H2O giving clear solution.
The infrared spectra were carried out on a Thermo‐
Fisher Nicolete FT‐IR spectrophotometer within range
1
(4000–400 cm−1). The H NMR spectra were measured
Fused heterocyclic systems containing pyrazole and
pyridine rings have been attracted considerable attention
due to their great usefulness in chemistry and a very wide
spectrum of their biological activities. Pyrazolopyridines
exhibit antitubercular, antiproliferative, inhibition of
cyclin‐dependent kinases and cardio‐vascular antiviral
and antileishmanial activities.[28–31] Pyrazolopyridines
specially, pyrazolo[3,4‐b]pyridine has proven to be inter-
esting class of heterocycles due to diverse biological prop-
erties including antioxidant and antitumor activities.[32]
This type of fused pryridines is isosters of bioactive
indoles or indazoles[33] representing important building
blocks in both natural and synthetic bioactive com-
pounds.[34] Various pyrazolo[3,4‐b]pyridines have been
shown antimicrobial, antiviral and antiparastic activi-
ties[35–38] and also exhibit pharmacological properties
such as anti‐inflammatory, anxiolytic activity along with
xanthine oxidase inhibitors, cholesterol formation inhibi-
tor and anti‐alzheimer. Also, pyrazolo[3,4‐b]pyridines
and their derivatives are attractive targets in organic syn-
thesis ascribed to their biological activities such as anal-
gesic, hypnotic, inhibitors of glycogen synthase kinase‐3
in DMSO‐d6 on a DELTA‐2 NMR spectrometer at
500 MHz, using tetramethylsilane (TMS) as internal refer-
ence. Fast Atom Bombardment (FAB) mass spectrum for
the ligand was performed on a Shimadzu Qp‐2010 Plus
spectrometer. The electron spin resonance (ESR) spectra
for Cu(II) complexes were performed on a Varian E‐
109c spectrometer equipped with a field modulation unit
at 100 kHz. The measurements were done in the X–band
with frequency 9.435 GHz on a microcrystalline powder
at room temperature; the microwave power was around
10 mW. Perkin‐Elmer Lambda 4B spectrophotometer
was employed to measure the absorption electronic spec-
tra in Nujol mulls. Molar conductivities of all metal com-
plexes (except Pd complex) were measured in DMSO
solution (10−3 M) at room temperature using a type
CD6N Tacussel conductimeter; while it is measured in
DMF solution for Pd(II) complex. The thermogravimetric
analyses (TG/DTG) were carried out using a Shimadzu
DAT/TG‐50 thermal analyzer with a heating rate of
10 °C/min under N2 atmosphere with a flowing rate of
20 mL/min from the room temperature up to 900 °C
using platinum crucibles. Powder X‐ray diffraction

EMAM ET AL.
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patterns were measured using an X‐ray diffractmeter
equipped with a graphite monochromator in the range
(2θ = 5°‐80°) by nickel separated CuKα radiation
(λ = 1.54060 A°).
Magnetic susceptibilities were measured at room tem-
perature by Gouy method using a Johnson Matthey mag-
netic susceptibility balance. The diamagnetic corrections
were calculated using Pascal's constants.[43] The effective
it was further stirred under reflux at 70 °C for 3 hrs. The
progress of the reaction was followed by TLC until all
starting substances were completely reacted. The formed
precipitate was filtered off, washed several times with hot
ethanol and dried in a vacuum desiccator over anhydrous
CaCl2. The melting point of the ligand was found 230 °C.
2.3 | Synthesis of Schiff Base‐metal
complexes
magnetic moments were calculated as Bohr Magneton
corrt
from the equation; μeff = 2.84(χM
T)1/2. The melting
points were measured with Stuart melting point appara-
tus. Biochemical measurements antioxidant and antibac-
terial activities) were evaluated at Biotechnology Center,
faculty of sciences, Mansoura University, Egypt.
Complexes (1–7) were synthesized by addition of
(10 mmol) of the appropriate metal salt dissolved in abso-
lute ethanol (15 ml) to a hot ethanolic solution (20 m)
containing (20 mmol) of HL ligand. The reaction mixture
was stirred under reflux for 5 hrs at 70 °C. The formed
complex was filtered off, washed several times by boiling
ethanol and dried under vacuum over anhydrous CaCl2.
2.2 | Synthesis of Schiff Base ligand; (Z)‐3‐
(1‐((4,6‐dimethyl‐1H‐pyrazolo[3,4‐b]
pyridin‐3‐yl)imino)ethyl)‐4‐hydroxy‐6‐
methyl‐2H‐pyran‐2‐one (HL)
2.4 | Biological activity
The starting material, 4,6‐dimethyl‐1H‐pyrazolo[3,4‐b]
pyridin‐3‐amine has been synthesized by the procedure
published previously[44] and also, the preparation of a
new ligand was carried out as shown in (Scheme 1). HL
ligand was prepared by refluxing equimolar ratios of
DHA (3.1 g) with 4,6‐dimethyl‐1H‐pyrazolo[3,4‐b]
pyridin‐3‐amine (3.0 g) in 40 ml of hot absolute ethanol
at 70 °C for 2 hrs. Then, few drops of DMF were added
to the reaction mixture as a catalyst (dehydrating agent);
2.4.1 | Antioxidant activity
The percentage of antioxidant activity (AA%) of HL
ligand and its metal chelates
was determined by 2,2‐diphenyl‐1‐picrylhydrazyl
(DPPH•) free radical assay. The DPPH• radical scaveng-
ing activity was measured according to spectrophotomet-
ric methodology.[45] As a result of the delocalization of
the DPPH• spare electron on the whole molecule, DPPH•
SCHEME 1 Synthesis of (HL) ligand

4 of 25
EMAM ET AL.
becomes a very stable free radical and not dimerize, as
happens with most free radicals. The delocalization on
the DPPH• molecule determines the occurrence of a deep
violet color, and absorption occurs at about 517 nm.
When DPPH• reacts with a hydrogen donor, the reduced
form is generated, accompanied with the disappearance
of the violet color. Therefore, the absorbance diminution
depends linearly on the antioxidant concentration.
Different concentrations of each investigated compounds
were prepared and taken in separate test tubes. The vol-
ume was adjusted to 100 μL with methanol. Five mL of
0.1 mM methanolic solution of DPPH• was added to
these tubes and shaken vigorously. The tubes were
allowed to stand for 20 min at 27 °C. The samples were
reacted with the stable DPPH• radical in methanol solu-
tion. The reaction mixture contains 0.5 ml of sample,
3 ml of methanol and 0.3 ml of DPPH• radical solution
(0.5 mM) in methanol. DPPH• radical is reduced forming
its molecule when it reacts with an antioxidant com-
pound, which can donate hydrogen. Mixture of methanol
(3.3 ml) and sample (0.5 ml) acts as blank. While, mixing
of 3.5 ml of methanol and 0.3 ml of DPPH• solution
affords the control solution and methanol was used for
the baseline correction. The experiment was performed
in triplicate for each compound and changes in the absor-
bance of the samples were measured.[46] After 100 min of
reaction, the changes in color from deep violet to pale
yellow were observed and read at 517 nm. Radical scav-
enging activity was expressed as the inhibition concentra-
tion (IC50), i.e. the concentration of extract necessary to
decrease the initial concentration of DPPH• by 50%
(IC50) under the specified experimental condition. The
scavenging activity percentage (AA%) was determined
according to the following equation[47]
Thereafter, 25 ml of nutrient agar (medium) was poured
into each petri plate and these plates were swabbed with
the test organism and kept 15 min for adsorption. They
were cultured at 37 °C for 18–24 hrs. The sterilized filter
paper discs (6 mm in diameter) were saturated with infu-
sion obtained from each sample and other set of filter
paper discs were soaked in DMSO as control. The discs
were placed on the surface of agar plates seeded with
the tested organism. The plates were incubated at 37 °C
for bacteria and diameters of inhibition zone (mm) were
measured after 18–24 hrs because they were taken as
measuring for antibacterial activity. Each treatment was
replicated three times. The antibacterial activity of a com-
mon standard antibiotic (Streptomycin) was also esti-
mated using the same procedure.
3 | RESULTS AND DISCUSSION
3.1 | Chemistry
The structure of HL ligand was identified using spectral
methods (mass, 1H NMR and infrared spectra). The Schiff
base ligands containing DHA can exist in different tauto-
meric forms as, di‐keto, keto‐enamine and enol‐imine
forms[49] as shown in (Figure 1). An intramolecular
hydrogen bridge binding one of the carbonyl groups
(C4=O) to the NH‐moiety of the imine group was
observed in (Figure 1b). Also, the intramolecular hydro-
gen bridge linking the azomethine group (C=N) to the
protonated OH moiety was found to be a characteristic
feature for HL ligand (Figure 1c and 1d). The six and four
membered intramolecular hydrogen bonding rings are
possible in the keto‐enamine and enol‐imine forms and
1
observed in (Figure 1). The IR and H NMR spectra con-
firmed the enol‐imine form (Figure 1c).
À
Á
Abssample − Absblank × 100
AA% ¼ 100 −
Where, A
Abscontrol
3.2 | Characterization of Schiff base ligand
(HL)
= absorbance of DPPH radical +
control
methanol and A sample = absorbance of DPPH• + sample
(test sample/standard).
3.2.1 | Mass Spectrum (FAB)
FAB mass spectral data of Schiff base ligand (HL)
(Figure S1) is carried out to restrict its molecular weight
and study its fragmentation pattern. The important frag-
ments obtained in its mass spectrum are shown in
(Scheme 2). The mass spectrum displayed a molecular
ion peak recorded at m/z = 332.18 peak corresponding
to [C16H16N4O3.H2O]+. The loss of one water molecule
giving a base peak at (m/z=313.21) which is equivalent
to molecular weight (312.33) and molecular ion
[C16H16N4O3]+. Further, this base peak loss variable rad-
icals affording fragment ion peaks recorded at m/z values
2.4.2 | Antibacterial activity
The ligand HL and its metal complexes were tested
in vitro antibacterial activity using the disc diffusion
method,[48] against E. Carotovora, P. Vulgaris (Gram –ve
bacteria); B. Subtillis, S. areus (Gram +ve bacteria). The
antibacterial activity of each sample was estimated by
paper disc diffusion assay using in inoculums containing
106 bacterial cells or 108 yeast cells/ml to be maintained
on a nutrient and Czapek Dox agar plates, respectively.

EMAM ET AL.
5 of 25
FIGURE 1 Tautomeric forms of HL
Schiff base
SCHEME 2 Proposed mass fragmentation of (HL) ligand
3.2.2 | 1H NMR spectra of HL ligand and
its Pd(II) complex
of 297.18; 269.17; 255.16; 228.17; 200.18; 187.17; 166.13;
151.11; 119.11; 104.1; 67.08 and 42.11 corresponding to
[C15H13N4O3]+;
[C13H9N4O3]+;
[C12H7N4O3]+;
[C11H6N3O3]+; [C9H4N3O3]+; [C9H3N2O3]+; [C7H3N2O3]
+; [C6H3N2O3]+; [C6HO3]+; [C6O2]+; [5.5C]+ and [3.5C]
1H NMR spectral data of HL ligand and its Pd (II) com-
1
plex are listed in (Table 1) and shown in Figure (S2). H
+
fragment ions, respectively.
NMR spectrum of Schiff base ligand in DMSO reveals

6 of 25
EMAM ET AL.
TABLE 1 1H NMR spectra of Schiff Base ligand and its palladium
(II) complex
In the 1H NMR spectrum of Pd(II) complex, the
absence of protonated OH indicates the release of the
OH proton and involvement of oxygen atom in binding
with the metal ion after deprotonation of OH group.
The up‐field shift of NH proton may be due to its involve-
ment in delocalization and its presence neighbor to nitro-
gen atom of cyclic (C=N)p.p group which is involved in
coordination, this is confirmed with IR spectra. The
methyl protons appeared at δ (2.49–2.61 ppm) on com-
plexation. In addition, the spectrum of complex depicts
a broad signal at δ (3.75 ppm) attributed to protons of
crystallized water.[56]
HL
[Pd(L)Cl].5H2O* Assignment
16.13
13.61
6.95
–
Enolic proton of –OH
12.80
6.83
6.06
2.61
NH proton
vinylic 14C‐H aromatic ring proton
5
5.87
vinylic C‐H aromatic ring proton
2.68
CH3 protons of azomethine
CH3 protons of P. P ring
2.53,2.56 2.53
2.15
2.49
CH3 protons of DHA moiety
(6C‐CH3)
3.75
Protons of crystallized water
3.2.3 | IR Spectrum of HL ligand
p.p: pyrazolopyridine ring,
1
*DMF is used in H NMR of PdCl2 complex.
The IR spectrum of HL ligand (Table 2 and Figure S3)
depicts most useful assignments for the characteristic
bands. The IR spectrum of the ligand pointed to the
absence of the bands characteristic to the amino group
and carbonyl of acetyl group and instead, a new
medium band due to azomethine (C=N) was appeared
singlet signals at δ (16.13 and 13.61 ppm) attributed to
protons of –OH and –NH, respectively.[50–52] It was found
that the protonated OH peak is deshielded relative to the
other OH protons, this is may be due to strong stabiliza-
tion of the enolic form by intramolecular hydrogen bond-
ing[53] (Figure 1c). While the singlet signals appeared at δ
6.95 and 5.87 ppm are denoted to two vinylic protons 14C‐
at1659 cm−1
.
This is indicative of the formation of Schiff base
ligand (Scheme 1). The spectrum of the ligand displays
spectral bands at (3451, 1156, 800); 1361 cm−1 cm−1 and
(1311, 1258, 1199) cm−1, assigned to protonated hydroxyl
group (ν(OH), δ(OH), γ(OH)); ν(C–N) of arylazomethine,
and phenolic ν(C–O), respectively.[14,21,23,56,57] These
vibrational bands are taken as evidence for the presence
of HL ligand in enol‐imine form (Figure 1c) rather than
5
H and C‐H, respectively.[51,52,54] Furthermore, the sig-
nals of protons of methyl groups attached to pyrane ring
and azomethine group appear as sharp singlet at δ 2.15
and 2.68 ppm, respectively. While, the signals of protons
of methyl groups attached with pyrazolopyridine moiety
observed at δ 2.52–2.56 ppm.[31,55]
TABLE 2 Fundamental infrared spectral bands (cm−1) of ligand (HL) and their assignments
Bands
Assignment
Bands
Assignment
3451(b, m)
υ(OH)enolic + υ (OH)H2O
υ(NH)P.P
1361(m)
υ(C‐N) arylazomethine + CH3 deformation
3220(w),3151(w)
3092(w)
3015(w)
υas(C‐H)aromatic
υs(C‐H)aromatic
1311(m),1258(w) 1199(m)
υ(C‐O)
enolic
2939(m),
2925(m)
2855(w)
2798(w)
1722(s)
υas(C‐H) of CH3
1156(m)
δ(OH)
800(m)
γ(OH)
υs(C‐H) of CH3
Hydrogen bonding
υ(C2 = O)lactone
υ(C=N)azomethine
1091(m), 945(s)
1065(m), 1030(w)
999 (w)
υ(N‐N)
(C‐H) in‐plane bending
υ(C‐O‐C)
γ(C2 = O)lactone
1659(m)
539(w)
1603(s)
1571(s)
υ(C=C)
709(w)
γ(NH) + (C‐H) out‐of‐plane bending
cyclic ring
υ(C=N)Py
υ(C=N)P.P
1497(m)
621(w)
505(w)
(C‐H) in‐plane bending of py
1462(w), 1442(w)
1402(w)
δas(C‐H) CH3
δs(C‐H) CH3
ρ(NH)P.P+(C–H) out‐of‐plane bending of py
Abbreviations: w: weak, m: medium, s: strong, b: broad, sh: shoulder, p.p: pyrazolopyridine, py: pyridine

EMAM ET AL.
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di‐keto and keto‐enamine forms (Figure 1a, 1b, 1d).[58]
The low value for protonated ν(OH) reveals its involve-
ment in strong intramolecular hydrogen bonding as
1
shown in Figure 1c. This is confirmed via its H NMR
spectrum.[59] Also, the IR spectrum (Table 2) indicates
to the medium bands at (3220, 3151) and 709 cm−1 attrib-
uted to stretching and bending vibrations of
pyrazolopyridine (NH) group.[54,60,61] On the other hand,
the spectrum displayed asymmetric and symmetric
stretching frequencies of aromatic and aliphatic (C–H)
which are characterized by 3015–3092; 2925–2970;
2855 cm−1, respectively.[53] But, the additional bands
observed at (1462, 1442); 1402; (1065, 1030) cm−1 are
restricted to asymmetric, symmetric bending vibrations
of (CH3) groups and aromatic (C–H) in‐plane deforma-
tion.[62] Moreover, the HL spectrum exhibits strong and
medium bands at (1722, 539); 1659; (1603,1571); 1497;
(999) and (945, 1091 cm−1) assignable to different vibra-
tions of lactone carbonyl (C=O); ν(C=N) of azomethine
group; (ν(C=C) + ν(C=N)) of pyridine; ν(C=N) of
pyrazolopyridine; ν(C–O–C) of pyrane ring and ν(N–N),
respectively.[49,56,63–65] Finally, the spectrum depicts weak
to medium bands at 621 and 505 cm−1 which are denoted
to (C–H) in‐plain and out‐of‐plane bending of pyridine
ring overlapped with ρ(NH) of pyrazolopyridine,
respectively.[53,54,66]
3.3 | Characterization of Schiff Base metal
complexes
3.3.1 | Analytical data
The physical and analytical results of HL ligand and its
metal complexes are tabulated in (Table 3). The structure
of the formed compounds was elucidated via microanaly-
ses, spectral studies, thermal (TG/DTG), magnetic and
molar conductance measurements. The metal complexes
obtained from the reaction of HL ligand with different
metal salts are separated with different stoichiometries.
All metal complexes are stable, non‐hydroscopic and par-
tially soluble in aprotic organic solvents, but they are
freely soluble in DMF and DMSO. The molar conductance
values are recorded within range; 10–48 Ω−1 cm2 mol−1
,
revealing to non‐electrolytic nature for complexes (1, 2,
4, 6 and 7) along with 1:1 electrolyte for complexes (3)
and (5).[67]
3.3.2 | IR spectra of metal complexes
FT‐IR spectroscopy is a very useful tool in the character-
ization of Schiff base metal complexes. The IR spectral
assignments of metal complexes are depicted in (Table 4

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EMAM ET AL.
and Figure S4). The spectra of all chelates display broad
via NNO type for complexes (1), (2), (4) and (6)), while, it
behaves as neutral bidentate species in chelates (3) and (7)
(NO fashion). This is supported by the existence of new
bands in the range (537–585) and (459–485 cm−1) attrib-
uted to ν(M–O) and ν(M–N),[55,58] respectively. In com-
plexes (3), (4) and (7), new absorption bands appeared
within range (3273–3446 cm−1), may be assigned to
ν(OH) of coordinated OH ion along with medium bands
observed at (1148–1150) and (803–807 cm−1) attributed
to δ(OH) and γ(OH) of coordinated OH.[71–73]
medium bands in the region of (3321–3550 cm−1) which
may be pointed to
ν
(OH) of lattice solvent
molecules/coordinated H2O.[63] Besides, in the spectra of
chelates (1–4) and (7), narrow weak bands at (779–
882 cm−1) specific for γ(H2O) rocking vibration of coordi-
nated water molecules are also seen.[68–70]
The bands ascribed to ν(NH) and γ(NH) of
pyrazolopyridine are displaced to lower/higher wavenum-
ber in all metal complexes due to the predominant tauto-
meric form of these complexes and the involvement of
adjacent (C=N) nitrogen atom in chelation with the metal
ion. The strong lactone carbonyl band (1722 cm−1) in the
free ligand is found to be unaltered on complexation
ruling out the participation of carbonyl oxygen in chela-
tion with the metallic centers, while this band is shifted
to lower wavenumber in complexes (1) and (6), this may
be due to the existence of strong inter and intramolecular
hydrogen bonding between carbonyl oxygen and hydro-
gen atom of lattice solvent molecules. Also, the ν(C=N)
band assignable to azomethine group exhibits blue/red
shift by (5–28 cm−1) and appeared in a weak nature at
(1631–1663 cm−1) leading to its involvement in chelation
via nitrogen atom in all complexes. This is supported by
upward/downward shift in ν(C‐N) values. While, the sec-
ond band characteristic to ν(C=N) of pyrazolopyridine is
positively and negatively shifted via (8–42 cm−1) and
(12 cm−1) in complexes (1), (2), (6) and (4), respectively.
Moreover, this band is unaltered in complexes (3), (5)
and (7). This indicates that the (C=N) of pyrazolopyridine
is one of the coordination sites in all complexes except
complexes (3), (5) and (7), this is confirmed through the
displacement of ligand ν(N–N) band (945, 1091 cm−1) to
higher/lower wavenumbers by (5–21), (10–35 cm−1) in
complexes (1), (2), (4) and (6).
The FT‐IR spectrum of copper(II) perchlorate com-
plex (5) revealed medium bands with high intensity
within range (3331–3453 cm−1) assigned to stretching
vibration of coordinated (OH) group which is masked
with that of lattice ethanol molecules. Moreover, the
δ(OH), γ(OH) and ν(C–O) bands exhibited blue/red shift
assisting the formation a bond between Cu(II) ion and
oxygen atom of phenolic group. Also, the free ligand band
(1659 cm−1) assigned to ν(C=N) of azomethine group is
splitted and shifted to lower wavenumbers (1653,
1631 cm−1) accompanied with weak nature in complex
(5), indicating the presence of two HL ligands chelated
with Cu(II) ion through azomethine nitrogen atoms. This
is supported by the red shift of arylazomethine (C–N),
(Table 4). On the other hand, the second (C=N) band of
pyrazolopyridine remains practically unaltered in fre-
quency and the change in ν(N‐N) is not observed, consid-
ering the noninvolvement of nitrogen atom of this group
in chelation. Also, the existence of perchlorate ion is
established by the appearance of strong band at
1089 cm−1 and medium band at 626 cm−1 assigned to
ʋ3(ClO4) and ʋ4(ClO4). The lack of splitting of these
bands show presence of non‐coordinated perchlorate.
Therefore, these values are matching with the molar con-
ductance values (Table 3), indicating the presence of ionic
perchlorate.[66,74–77]
The IR spectra of metal complexes (Table 4) display
that the protonated ν(OH) is not clearly detected due to
its overlapping with ν(OH) of lattice/coordinated solvent
molecules. So, it is difficult to detect the shift in this band.
While, the in‐plane and out of plane bending vibrations
(δ(OH), γ(OH)) bands of protonated OH appeared at
(1156, 800 cm−1) in the free ligand (Table 2) disappeared
in complexes (1), (2), (4), (6), and the bands appeared at
(1199, 1258, 1311 cm−1)assignable to ν(C–O) exhibit alter-
ation in its shape and position with different degrees.[14]
These observations indicate the involvement of phenolic
(C–O) in coordination after deportation.[64] However,
the free ligand bands (δ(OH) and γ(OH)) exhibit blue shift
via (6–8 cm−1) and red/blue shift via (5–22 cm−1) in metal
chelates (3) and (7). This behavior indicates the chelation
of HL ligand through phenolic (–OH) group oxygen with
Cu(II) and Fe(III) ions. These results reveal that HL
ligand chelates with ions in monobasic tridentate manner
The aforesaid changes in the IR spectral bands
established that HL ligand behaves as a monobasic
tridentate fashion through (NNO) type in complexes (1),
(2), (4) and (6). While, HL chelated as (NO) type in com-
plexes (3) and (7) with the metallic ions through nitrogen
and oxygen atoms of azomethine group and phenolic –
OH, respectively. But, the two ligands behaved as mono-
basic and neutral bidentate fashions via (NO) type in
complex (5).
3.3.3 | Electronic spectra and magnetic
moment measurements
The electronic spectra and magnetic susceptibility mea-
surements have been used to determine the possible
geometry of the prepared metal complexes. The electronic

EMAM ET AL.
11 of 25
spectra of HL ligand and its metal complexes were mea-
sured in Nujol mull at room temperature. The magnetic
moment values and electronic absorption data are listed
in (Table 5).
The spectrum of the ligand displays three bands, at
256, 301, 374 and 456 nm. The first, second and third
bands are due to the electronic transition between
anti‐bonding orbitals (π–π*) of heterocyclic moiety,
aromatic rings along with carbonyl and (C═N)
groups.[58,78] The fourth band attributable to n–π* transi-
tion of conjugated system of the ligand.[79] The spectra of
all metal complexes showed bands at 242–290 and 300–
321 nm due to π–π*, while the electronic transition bands
of n–π* show up at 422–438 nm.[80] The π–π* and n–π*
transitions of the ligand suffered low shift upon complex-
ation. All complexes display bands at (424–498 nm)
TABLE 5 Electronic spectral data and magnetic moments of the ligand and its metal complexes
μeff. (B.M.)
No.
Compound
Spectral bands(nm)
Assignment
(per metal ion)
HL. H2O
256, 301
374
π‐π*
–
456
n‐π*
1
[Co(L)Cl(H2O)].0.5EtOH.3.75H2O
[Cu(L)Cl(H2O)2].EtOH.2.5H2O
246, 287
311,432
470
498
632
Intraligand transitions
4.80
LMCT
4
4A2’(F) → A2’(P)
4
4
́
A2’(F) → E”(P)
4
4
́
A2’(F) → E'(F)
695
2
242,316
368,432
468
528
572
Intraligand transitions
1.79
LMCT
2
2B1g → Eg
2B1g → B2g
2B1g → A1g
2
2
670
3
4
[Cu2(HL)3Br2(OH)(H2O)]Br.6.5H2O
[Cu(L)(OH)(H2O)].0.25EtOH.0.5H2O
246, 289
336,438
498
552
690
Intraligand transitions
1.12
1.80
2
2B1g → Eg + LMCT
2
2B1g → B2g
2
2B1g → A1g
246,268
300,355
466
540
721
Intraligand transitions
LMCT
2
2B1g → Eg
2B1g → A1g
2B1g → B2g
2
2
769
5
[Cu (HL)(L)]ClO4.4.5EtOH
248,290
321,422
466
670
714
Intraligand transitions
1.72
LMCT
2
2B1g → A1g
2B1g → Eg
2
2
786
2B1g → B2g
6
7
[Pd(L)Cl].5H2O
248,289
311
424
468
509
Intraligand transitions
Diamagnetic
1
LMCT+1A1g → E1g
1
1A1g → B1g
1
1A1g → A2g
[Fe (HL)Cl (OH)2(H2O)].0.75EtOH.1.25H2O
252,288
316,422
450
476
542
Intraligand transition
4.85
LMCT
4
6A1g(S) → A1g(G),4Eg(G)
6A1g(S) → T2g(G)
6A1g(S) → T1g(G)
4
4
681

12 of 25
EMAM ET AL.
range, assignable to ligand to metal charge transfer
(LMCT).[80]
3.4 | Thermal studies
Thermal analysis is a very useful technique in order to
restrict the thermal stability of compounds, confirm
molecular formulae and decide the different types of sol-
vent of crystallization. This measurement is used to sug-
In fact the μeff values for trigonal bipyramidal cobalt
(II) complexes lie in the range 4.26–5.03 BM, owing to
the orbital contribution of the excited (E') levels. The
effective magnetic moment of (1) is (4.80 BM) as in some
of the reported high‐spin cobalt(II) complexes. This is
because the orbital contribution to the spin‐only value
(3.87 BM) is fairly small, may be due to the presence of
a low symmetry ligand field component with tetrahe-
drally distortion.[81,82] Cobalt(II) complex (1) displays
gest
a
general mechanism for the thermal
decomposition of these complexes. Figure (S5) represents
TG/DTG curves of HL ligand and its complexes (1–7).
The TG curve of complexes reveal mainly three decompo-
sition steps except complexes (1), (4) and (7) which
undergo thermal degradation in four stages. The decom-
position of all metal complexes initiated at low tempera-
ture (21–30 °C), this is may be due to the physically
attached of the solvent of crystallization with the coordi-
nation sphere through inter/intramolecular hydrogen
bonding. The final residue calculated from thermal
decomposition is metal for complexes (2–4 and 6) and
metal oxide contaminated with carbon for complexes (1,
5 and 7). The structural properties of HL ligand, metal
chelates, the different types of anions and the variety of
the metal ions influence on the thermal behavior of the
desolving/decomposition step in metal complexes, this
studying can be explained by calculating the activation
energy (E*), the order of reaction (n), the activation
enthalpy (ΔH*), the activation entropy (ΔS*) and the free
energy of activation (ΔG*) using the analysis of the DTG
curves. The order of reaction (n) can be determined from
electronic absorption bands at 498, 632 and 695 nm,
4
assigned to 4A2’(F) → 4A2’(P), A2’(F) → 4E”(P) and
́
A2’(F) → 4E'(F) transitions, respectively. These bands
4
́
suggest a trigonal bipyramidal geometry for Co(II)
complex.[72,81,82]
The spectra of Cu(II) complexes (2) and (3) exhibit
bands at (498–528), (552–572) and (670–690) nm assigned
2
2
2
2
2
2
to B1g → Eg, B1g → B2g and B1g → A1g transitions,
respectively. These bands are consistent with octahedral
geometry.[60,72,74,83] But, Cu(II) complex (4) shows bands
2
2
at 540, 721 and 769 nm, corresponding to B1g → Eg,
2B1g → 2A1g and 2B1g → 2B2g, indicating to square pyrami-
dal structure for complex (4).[84–86] Also, the electronic
spectrum of complex (5) depicts bands at 670, 714 and
786 attributed to 2B1g → 2A1g, 2B1g → 2Eg and 2B1g → 2B2g
transitions affording square planar geometry for
Copper(II) ion. The magnetic moment values of Copper(II)
complexes are in the range (1.12–1.80 B.M) expected
for distorted tetragonal structure around Cu(II)
ion.[27,54,60,87] The lower magnetic moment value (1.12 B.
M) for complex (3) pointes to the existence of interaction
between Cu(II) ions, this is evidenced by the exchange cou-
pling parameter (G) in its ESR spectrum.[66,74]
the relation: (1–αm) = n(1–n) where (αm) is the
,
decomposed fraction at DTG peak (the maximum rate
of decomposition). The activation energy values of (E*)
for the desolvation/decomposition stage can be calculated
using the initial rate method,[91,92] where, (lnI) is plotted
versus (1000/T) for the low‐temperature peak affording
a straight line with a slop equals to (−E*/R), where (I)
represents (dα/dt) in DTG curve and ΔT in DTA
curve.[73,92,93] Also, the parameters (ΔH, ΔS* and ΔG)
were estimated from these relations: ΔH = E*–RΔT;
ΔS* = RTln (Ah/KT); ΔG = ΔH–TΔS*, where K, h, R, A
and T are Boltzmann's, Planck's and gas constants, fre-
quency factor and absolute temperature, respectively.
The thermal decomposition pathway of HL ligand and
its metal complexes was studied by using TG analysis.
The TG and DTG curves (Figure S5) together with the
thermal results and their assignments are listed in
(Table 6). So, the thermodynamic and kinetic parameters
of the metal complexes are given in Table 7. The TG mea-
surements were performed at a heating rate 10 °C/min
under nitrogen within temperature range 21–900 °C.
The TG curve of ligand indicated that its thermal
decomposition achieved through three steps. The first
step takes place at temperature range 31–230 °C referring
to small weight loss by 2.73% due to removal of 0.5 mol of
Diamagnetic Pd(II) complex displays three bands at
424, 468 and 509 nm assignable to 1A1g → 1Eg, 1A1g → 1B1g
1
1
and A1g → A2g, respectively, which is compatible with
square planar stereochemistry.[72,74,87,88]
Iron(III) ion is usually high‐spin in nearly all of its
complexes, except those containing strongest ligands.
The high‐spin d5 configuration of octahedral Fe(III) is
expected to afford a magnetic moment, nearly equal to
the spin‐only value (5.90 B.M), where, the ground state
has no orbital angular momentum. Lower values are
reported for some hydroxo, chloro and bromo Fe(III)
complexes due to the spin‐crossover phenomenon.[89] It
is to be noted that the magnetic moment value for com-
plex (7) is (4.85 BM) implying high‐spin octahedral struc-
ture around Fe(III) ion. The spectral bands of Fe(III)
complex appear at 476, 542 and 681 nm corresponding
6
4
6
4
to A1g(S) → A1g(G),4Eg(G), A1g(S) → T2g(G) and
6A1g(S) → 4T1g(G) transitions indicative of high‐spin octa-
hedral geometry.[83,90]

EMAM ET AL.
13 of 25
TABLE 6 Thermal analyses (TG/DTG) data for Schiff Base ligand (HL) and its metal complexes
Mass Loss%
TG
DTG
No. Compound
range (°C) Peak (°C)
Calcd. Found Assignment
Ts (°C)
HL.H2O
31–230
2.73
2.73
Loss of 0.5 mol of crystallized
H2Oa
230
230–280
261(s)
20.93 21.59
Loss of 0.5 mol of lattice H2O
and partial ligand pyrolysis
(C4H12)c
280–900
At 900
69.07 67.97
ligand pyrolysis (C10H4N4O3)c
(2C)d
7.27
3.11
7.71
2.95
1
[Co(L)Cl(H2O)].0.5EtOH.3.75H2O
24–110
49(w)
Loss of 0.25 mol of EtOH and
0.25 mol of crystallized
200
210
240
1+2
H2O
110–200
200–330
163(br.w)
229(br)
12.75 12.56
12.16 12.40
Loss of 0.25 mol of EtOH, 3 mol
of crystallized H2O
1+2
Loss of 0.5 mol of H2O, one mol
of coordinated H2O and
0.5 mol of Cl2 gasc
Partial ligand pyrolysis
(C5H4O)c
Complete ligand pyrolysis
(C8H11N4)c
(Co2.54O4+ 3C)d
330–454
454–750
374(w)
15.57 16.26
31.73 31.28
24.68 24.53
580(s.br)
At 750
2
[Cu(L)Cl(H2O)2].EtOH.2.5H2O
21–114
49(w)
3.81
3.18
Loss of 0.25 mol of EtOH and
0.5 mol of crystallized
1+2
H2O
114–210
210–310
310–776
At 776
170(m)
270(w)
697(m.br)
10.62 10.98
12.47 11.99
61.28 61.27
11.81 12.50
Loss of 0.75 mol of EtOH and
1.25 mol of crystallized
1+2
H2O
Loss of 0.75 mol of lattice H2O,
one mol of coordinated H2O
and 0.5 mol of Cl2 gasc
Complete ligand pyrolysis and
loss of one mol of coordinated
H2Oc
(Cu)d
3
[Cu2(HL)3Br2(OH)(H2O)]Br.6.5H2O
21–85
Stable zone
85–176
142(w)
5.88
2.17
6.23
2.25
Loss of 4.75 mol of crystallized
H2O moleculesa
Loss of 1.75 mol of crystallized
H2O moleculesa
176–240
240–361
361–850
270(m)
733(s.br)
10.98 10.69
72 25 72.17
Loss of one mol of Br2 gasc
Loss of one mol of coordinated
H2O, one mol of OH, 0.5 mol
of coordinated Br2 gas and
3 mol of HL ligandc
At 850
8.72
2.68
8.71
2.39
(2Cu)d
4
[Cu(L)(OH)(H2O)].0.25EtOH.0.5H2O
25–185
185–304
47(w)
Loss of 0.25 mol of EtOHb
Loss of one mol of coordinated
H2O and partial ligand
pyrolysis (C5H4O)c
185
253(m)280(s) 22.80 22.43
304–410
380(s.br)
24.89 24.65
Loss of 0.5 mol of lattice H2O,
one mol of coordinated OH
and partial ligand pyrolysis
(C5H7N)c
(Continues)

14 of 25
EMAM ET AL.
TABLE 6 (Continued)
Mass Loss%
TG
DTG
No. Compound
range (°C) Peak (°C)
Calcd. Found Assignment
Ts (°C)
410–900
592(w)
34.88 35.78
Complete ligand pyrolysis
(C6H4N3O2)c
(Cu)d
At 900
14.75 14.76
5
[Cu (HL)(L)]ClO4.4.5EtOH
23–194
194–246
246–371
Stable zone
246
–
2.32
2.02
Loss of 0.5 mol of EtOHb
274(s)
22.76 22.22
31.32 30.89
33.19 34.54
10.41 10.35
Loss of 2.75 mol of lattice EtOH
and one mol of ClO4
Loss of one mol of ligand
(C16H15N4O3)c
Loss of 1.25 mol of EtOH and
ligand pyrolysisc
(CuO+ 2C)d
c
371–534
534–700
507 (m,br)
572(m,br)
At 700
6
[Pd(L)Cl].5H2O
24–90
37(w)
3.32
4.14
9.12
3.10
4.39
9.17
Loss of one mol of crystallized 280
H2Oa
90–237
Loss of 1.25 mol of crystallized
H2Oa
237–280
280–550
268(w)
409(m,br)
Loss of 2.75 mol of lattice H2Oa
Loss of 0.5 mol of Cl2 gas and
ligand pyrolysisc
63.83 63.72
At 550
19.59 19.64
(Pd)d
7
[Fe (HL)Cl (OH)2(H2O)].0.75EtOH.1.25H2O 25–104
44(w)
277(s)
2.24
4.49
19.71 19.01
2.26
4.38
Loss of 0.25 mol of EtOHb
Loss of 0.5 mol of EtOHb
Loss of 0.75 mol of lattice H2O,
one mol of coordinated H2O,
two mol of coordinated OH
and 0.5 mol of Cl2 gasc
235
104–235
235–300
300–356
323(m,br)
400(s)
15.23 15.60
Loss of 0.5 mol of lattice H2O
and partial ligand pyrolysis
(C4H7N)c
356–600
30.46 30.64
27.87 28.12
ligand pyrolysis
(C6.75H9N3O1.5
c
)
At 600
(0.5Fe2O3+ 5.25C)d
Abbreviations:
a: dehydration
b:desolvation
c:decomposition,
d: final residue, TS: thermal stability, s: strong; m: medium; w: weak and br: broad
crystal water molecule. The ligand starts its decomposi-
tion in the second step at 230–280 °C by weight loss
21.59% assigned to elimination of the rest of water mole-
cule and partial ligand pyrolysis (C4H12). This process
associated with strong DTG peak at 261 °C. Then, this
step is followed by complete degradation at temperature
range 280–900 °C with weight loss of 67.97% leaving car-
bon residue with 7.71%.
The TG curves of complexes (2), (3) and (6) reveal that
they decompose by similar way. The initial desolvation
step is observed through two processes at temperature
ranges, 21–114, 114–210; 85–176, 176–240 and 24–90,
90–237 °C associated with weak to medium DTG peaks
at 49,170; 142 and 37 °C for complexes (2), (3), and (6),
respectively, (Scheme 3). Hence, the low temperature
range for this transformation is indicative of the existence
of releasing of adsorbed lattice solvent in chelated com-
pounds. In this process, the ease of desolvation for these
solvent molecules is ascribed to weak interaction in the
crystal voids or no role in the lattice forces.[66,74,87,88,94]
The calculated activation energies are found to be 75.99,
173.1; 75.91 and 140.72 KJmol−1 for complexes (2), (3)
and (6), respectively. After that, the second decomposi-
tion step observed at temperature ranges (210–310),
(240–361) and (237–300 °C) is accompanied by
(Tmax = 270, 270, 268 °C) with mass losses 11.99, 10.69

EMAM ET AL.
15 of 25

16 of 25
EMAM ET AL.
and hydroxide) for complex (3) and coordinated chloride
ion for complex (6). This step is accompanied with strong
broad DTG peaks at (Tmax = 697, 733,490 °C), ended with
the formation of copper and palladium metals for com-
plexes (2), (3) and (6), respectively.[72,74,88]
The thermoanalytical data showed that complexes (5)
and (7) have isothermal behavior. The thermal decompo-
sition proceeds in four steps within temperature range
(194–700 °C) and (25–600 °C), where, Cu(II) complex
shows no weight up to 194 °C. The first one (Schemes 4,5)
occurs through one or two events within the temperature
ranges (194–246 °C) and (25–235 °C) with weight losses
2.02% and 6.64%. This is pointing to partial desolvation
with the removal of 0.5 and 0.75 mol of crystallized sol-
vent for complexes (5) and (7), respectively. The thermal
stability is observed up to 246 and 235 °C for complexes
(5) and (7), respectively. After partial desolvation, the sec-
ond decomposition stage takes place up to 371 and 300 °C
assignable to removal of lattice EtOH molecules along
with ionic perchlorate ion for complex (5) and elimina-
tion of lattice water, coordinated water molecule, chloride
ion as well as coordinated hydroxyl for complex (7). This
step is accompanied with sharp strong DTG peaks at
(Tmax = 274 and 277 °C) and the estimated activation
energy for this step was found (388.18 and 474.95
KJmol−1). The last degradation steps (third and fourth)
in complexes (5) and (7) refer to loss of organic moieties
at temperature ranges (371–700) and (300–600 °C)
together with the remain lattice solvent molecules giving
CuO and 0.5 Fe2O3 mixed with carbon.[54,57,66,72]
The thermal decomposition of complexes (1) and (4)
shows four steps of releasing weight within temperature
ranges (24–200), (185–330), (304–454) and (401–900 °C).
The first decomposition step at (24–200 °C) is ascribed
to partial desolvation with weight loss of 15.51 and
2.39% associated with weak DTG peak at (Tmax = 49,
163 and 47 °C), respectively (Table 6; Scheme 6). On
the other hand, this step is represented by two events
with activation energy 88.59, 24.86 KJmol−1 and order
of reaction 0.9, 1.58 for complex (1), while, this step is
shown by single event with activation energy 70.40
KJmol−1 and first order reaction for complex (4)
(Table 7). The complexes (1) and (4) begin the second
degradation process at 200 and 185 °C up to 330 and
304 °C, respectively. This step is observed as one event
with a broad DTG peak at (Tmax = 229 °C) for complex
(1) and two successive events with strong DTG peaks at
(Tmax = 253,280 °C) for complex (4). This was assigned
to elimination of 0.5 mol of lattice H2O molecule, one
mol of coordinated H2O molecule and 0.5 mol of Cl2
gas for complex (1), in addition to partial ligand pyroly-
sis with coordinated H2O in case of complex (4)
(Table 6). When the temperature increases, within
SCHEME 3 Thermal decomposition mechanism of complexes
(2) and (6)
and 9.17% is assigned to removal of 0.75 mol of lattice
H2O molecule, 0.5 mol of Cl2 gas along with removal of
coordinated H2O molecule for complex (2). The counted
activation energies for this step are within range
(121.68–336.37 KJmol−1). But, this step corresponds to
elimination of 0.5 mol of Br2 gas and 2.75 mol of lattice
water molecules for complexes (3) and (6), respectively.
Finally, this step is followed by complete ligand pyrolysis
for all complexes together with releasing of coordinated
H2O for complexes (2,3); coordinated anions (bromide

EMAM ET AL.
17 of 25
SCHEME 4 Thermal decomposition mechanism of complex (5)
SCHEME 5 Thermal decomposition mechanism of complex (7)

18 of 25
EMAM ET AL.
SCHEME 6 Thermal decomposition
mechanism of complexes (1) and (4)
temperature ranges (330–750 °C), complete ligand pyrol-
ysis takes place ending with the formation of cobalt
oxide contaminated with carbon for complex (1),
whereas, the removal of coordinated and lattice water
along with complete delegation occurs for complex (4)
within (410–900 °C) rang, and ended with copper metal
as final residue[72,74,88] (Scheme 6, Table 6). The struc-
ture of cobalt oxide residue was illustrated using XRD
diffraction pattern (Figure 2). It shows that no impurity
peaks are observed, implying that the formed residue
was pure Co2.54O4 and it is consistent with the card
number (JCPDS 01–078‐5623).[95] XRD pattern reveals
all peaks intensities of (531), (440), (511), (400), (222),
(311), (220) and (111), which can be perfectly indexed
to diffraction angles (2θ in degree); 77.01, 64.95, 59.08,
44.48, 38.24, 36.50, 30.99 and 18.92, respectively. The
crystallite size of the investigated Co2.54O4 oxide was
calculated from the line broadening analysis of some
diffraction
lines
utilizing
Scherer
equation[96]
;
(D = Kλ/βCosθ), where (K), (D), (λ), (β) and (θ) are
Scherer's constant (0.90), the mean crystallite size
(nm), the X‐Ray wavelength (A°), full width half maxi-
mum (FWHM) of the oxide diffraction peaks and dif-
fraction angle. The average crystallite size of this oxide
was counted (39.05 nm).
3.5 | Kinetic and thermodynamic studies
Based on TG/DTG curves and the calculated thermody-
namic parameters, the investigation of Tables 6 and 7
established these remarks:

EMAM ET AL.
19 of 25
between Cu(II) ion with N2O2 chromophore in square
planar Cu(II) complex (5) along with the increase in
chelate rings around the central ion.[66,72]
• Table 7 demonstrates that the activation energy of
desolvation step for complex (2) (75.99KJmol−1) is
greater than that observed for complex (4)
(70.40KJmol−1). This is attributed to the number and
the nature of crystallized solvent molecules[68,87]
which is confirmed with the higher thermal stability
of complex (2) (Ts = 210 °C) than that of complex
(4) (185 °C) (Table 6).
• The negative entropy values (ΔS*) of complexes for the
thermal decompositions steps (Table 7), showed that
the transition states are more ordered and the thermal
reactions proceeded in complicated mechanisms.
FIGURE 2 XRD pattern for cobalt oxide residue isolated from
TG decomposition of complex (1)
• The Cu(II) perchlorate complex (5) has a higher thermal
stability (start temperature of decomposition) than
Cu(II) bromide complex (3). This may be due to the dis-
tribution of the ligand molecules around the central
metal ion as shown in Figure 3, in addition to the differ-
ent geometrical structures.[74] It was found that complex
(3) has more steric effect than that observed for complex
(5) (Figure 3). Also, the lower thermal stability of com-
plex (3) is attributed to the strong repulsion between
electrons of the two Cu(II) ions which has weakened
on M–Br bond and facilitate the rupture of four mem-
bered ring in the first step decomposition reaction due
to its exposed to torsional strain. This was confirmed
by the higher activation energy value for the initial
decomposition of complex (5) (388.18KJmol−1) than
that required for complex (3) (202.94 KJmol−1).
3.6 | X‐ray diffraction studies of metal
complexes
The X‐ray diffraction patterns of the prepared metal che-
lates (3), (4) and (6) are observed in curves a, b, and c of
Figure 4, respectively. It is observed that Cu(II) bromide
(3) and Cu(II) sulphate (4) complexes have polycrystal-
line nature. The X‐Ray diffraction pattern of complex
(3) is consistent with the peak positions in the (JCPDS
cards, No. 00–064‐1623, 01–082‐2118, 01–076‐5674) show-
ing diffraction angles at 9.59, 11.47, 14.43, 16.40, 18.59,
23.46, 25.81, 27.09, 32.80, 36.73°. The low‐angle reflection
peaks at 11.47, 16.40, 25.81° can be found with d‐spacing
of 7.71, 5.37, 3.45 A°, which perfectly matched the
reported interlayer spacing value for (C2H4(OH)
NH3CuBr2). Whereas, the broad peak with very low
intensity at around 27.09° was assigned to be the traces
of CuBr with Cu(II) bromide. In addition, the weak inten-
sity peaks at 18.59 and 25.81° matched with C9H14O7 for-
mula. The calculated crystallite size of complex (3) using
Scherer equation was found (20.83 nm). Also, the XRD
curve of complex (4) is consistent with the reported
(JCPDS card, No. 01–073‐4721), showing strong and weak
diffraction peaks at 8.56, 9.35, 11.67, 14.32, 18.97, 23.33,
27.04, 28.72, 42.20 and 40.72° which perfectly matched
with formula (O2CN(CH3)2)2(NH (CH3)2)2Cu. The aver-
age crystallite size of this complex (4) exhibited a smaller
average crystallite size of about 12.87 nm, than that of
complex (3). While, Pd(II) complex (6) can have amor-
phous nature.
• The higher thermal stability of square planar Cu(II)
complex (5) (246 °C) than the distorted octahedral
Cu(II) complex (2) (210 °C) may be due to the ionic
character of complex (5) and the stronger interaction
3.7 | Electron spin resonance studies
The ESR spectra of Copper(II) complexes (2–5) in the
polycrystalline state were scanned at room temperature
and the geometrical parameters were calculated. The data
FIGURE 3 Distribution of chelate rings around Cu (II) ion in
complexes (3) and (5)

20 of 25
EMAM ET AL.
TABLE 8 ESR spectral parameters of Copper(II) complexes of
HL ligand
Complex
Parameter (2)
Complex
(3)
Complex
(4)
Complex
(5)
g∥
–
2.870
2.199
2.051
2.10
2.530
g⊥
–
2.130
2.094
agav/giso
bG
2.12
–
2.380
2.240
3.10
–
4.45
5.75
–
cAǁ
cA⊥
dAiso
Kǁ
–
147.70
63.83
91.79
0.621
0.738
0.699
0.668
0.578
0.816
148.89
13004
18519
–
–
–
–
–
–
–
–
–
K⊥
eK
α2
β2
γ2
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
f f
–
–
–
gΔ1
gΔ2
–
–
–
–
–
–
gǁ þ 2g⊥
a
b
c
gav
=
3:
gǁ − g
G =
e , ge= 2.0023.
g⊥ − ge
= unit of A in (cm−1).
Aǁ þ 2A⊥
d
Aiso
=
3
Kǁ þ 2K⊥
3
e
f
K =
gǁ
Aǁ
f =
in (cm).
gΔ1= 2B1g→2B2g (in cm−1).
gΔ2= B1g →2Eg (in cm−1).
2
FIGURE 4 XRD curves for complexes (3), (4) and (6)
Copper(II) ion in addition to the interaction between
the unpaired electron and the magnetic nuclei.[97]
Copper(II) complexes (3, 4, 5) depict axial spectra with
two (g) values (g|| = 2.199–2.87, g⊥ = 2.051–2.13) associated
were listed in (Table 8) as well as the ESR curves for com-
plexes (2) and (4) which are shown in Figure (S6). It was
known that the Copper(II) ion have d9 configuration with
effective spin of (s = 1/2) and spin angular momentum of
(ms = 1/2) creates a doubly degenerate spin state in the
absence of external magnetic field. While the degeneracy
of this state is lifted in the presence of a magnetic field
and splitted to two spin energy levels with energy differ-
ence which is given by: ΔE = hυ = gβH, where (h) is
Planck's constant, (υ) is the microwave frequency, (β) is
Bohr magneton, (H) is the applied magnetic field and
(g) is the land splitting factor for electronic transition
from ms = +1/2 to the other level of ms = −1/2.
2
with the trend g|| > g⊥ > ge indicating of B1g as ground
state.[87,98] This is agreement with the electronic configura-
tions (eg)4(b2g)2 (a1g)2 (b1g)1; (eg)4(a1g)2(b2g)2(b1g)1 and
(a1g)2(eg)4(b2g)2(b1g)1 which are assignable to distorted octa-
hedral, distorted square pyramidal and square planar geom-
etry for complexes (3), (4) and (5), respectively.[98] In these
chelates, the lowest (g)value is greater than (2.04)
(Table 8), indicating to axial symmetry and all principle
axes are aligned parallel with a (dx2‐y2) ground state for
square pyramidal geometry.[99–101]
The geometric parameter, (G=g||‐2/g⊥‐2), for this type of
symmetry determines the exchange interaction between
Cu(II) ions in the polycrystalline form. The observed value
of this complex is equal to (4.45–5.75) suggesting the
absence of copper‐copper interaction in complexes (4) and
(5).[87,102] Whereas, the calculated G value for complex (3)
The ESR spectrum for complex (2) displays isotropic
spectrum accompanied with (g) value of 2.12 and two
broad signals with no hyperfine structure. It was
concluded that the broadness of this peak is ascribed to
dipolar interaction and
a random orientation of

EMAM ET AL.
21 of 25
was found (3.10) indicative of the existence of Cu‐Cu inter-
action, this is in agreement with the lower value for its mag-
netic moment (μeff = 1.12BM). The ESR spectrum of
complex (4) shows hyperfine splitting due to interaction
between nuclear spin moment of Cu+2 ion (I = 3/2) and
its electron spin moment (s = 1/2) in the parallel region.
On the other hand a weak hyperfine splitting is observed
in the high field region (Table 8). Also, the superhyperfine
splitting due to coordinated nitrogen or oxygen is not shown
for complex (4). The g||, A||, g⊥ and A⊥ values are listed in
(Table 8). On the other hand, the empirical factor f
(f = g||/A||) can be utilized to determine the stereochemistry
of Cu(II) complexes. The (f) value is counted (148.89 cm)
for complex (4) pointing to tetragonal distorted Copper(II)
complexes.[74,92,103,104] In addition, the nature of metal–
ligand bond is characterized by using the energies of d‐d
electronic transitions. The bonding parameters α2, β2 and
γ2 which are taken as measure of covalence of in‐plane σ
bonds, in‐plane and out‐of‐plane π bonds; respectively will
be calculated by these equations[90,104–106]
:
ꢀ
ꢁ
A‖
3ðg⊥ − geÞ
∝2 ¼
þ g∣ − ge
þ
þ 0:04
0:036
7
ðg‖ − geÞΔ1
8λ0
∝2β2 ¼ K2‖ ¼
ðg⊥ − geÞΔ2
¼
2λ0
α2γ2 ¼ K2⊥
Where λ° is the spin‐orbit coupling constant for Cu(II),
(α2) indicates to the strength of (σ) bond between the
ligand and Cu(II) ion, (K||, K⊥) are parallel and perpendic-
ular components of orbital reduction factor and (Δ1,Δ2) are
2B1g → B2g and B1g → Eg, respectively. In our studied
complex (4) (α2 = 0.668) indicating that the metal–ligand
(σ) bond has slightly covalent character.[107]According to
Hathaway,[108] K < 1 and K|| < K⊥ assigned to in‐plane π‐
bonding as well as greater covalent nature for metal–
ligand bond. Also, the β2 and γ2 coefficients are less than
unit attributed to the presence of in‐plane and out‐of‐plane
π character for metal–ligand bonds.[74]
2
2
2
FIGURE 5 Proposed structures for metal complexes of (HL)
ligand
low concentrations compared to those of an oxidizable
substrate. The oxidation reaction produces free radicals
implying to chain reactions which destroy cells. So, anti-
oxidant prevents the initiation or propagation of oxidative
chain reactions. In addition, the presence of antioxidants
in food is very important for health‐protection because
they minimize the risk for chronic diseases including can-
cer and heart disease. The activity of a compound on the
DPPH• radical is a fast and reliable parameter to measure
the in vitro antioxidant activity of sample. This assay is
based on the measurement of the decrease in molar
Based on elemental analysis, spectral methods and
thermal studies, the proposed structures for metal com-
plexes are shown in Figure 5.
3.8 | Biological activity
3.8.1 | Antioxidant activity
Antioxidants are molecules which inhibit or delay the
oxidation of other molecules when they are found at

22 of 25
EMAM ET AL.
TABLE 9 In vitro antioxidant assay of HL ligand and its metal complexes using DPPH• method
Inhibition %
Concentration mg/mL
No.
Compound
IC50
0.009
0.004
HL
0.0103
0.235
0.0211
0.156
0.076
0.078
0.1337
0.103
0.0221
73.36
51.3
65.95
52.80
55.9
56.90
45.43
50.5
–
72.29
50.61
60.23
–
1
2
3
4
5
6
7
[Co(L)Cl(H2O(].0.5EtOH.3.75H2O
[Cu(L)Cl(H2O)2].EtOH.2.5H2O
[Cu2(HL)3Br2(OH)(H2O)]Br.6.5H2O
[Cu(L)(OH)(H2O)].0.25EtOH.0.5H2O
[Cu (HL)(L)]ClO4.4.5EtOH
[Pd(L)Cl].5H2O
47.63
50.99
–
[Fe (HL)Cl (OH)2(H2O)].0.75EtOH.1.25H2
Ascorbic acid
46.34
–
All the concentrations in mg/ml at concentration 4.6 mg/ml
TABLE 10 In vitro antibacterial activity of HL ligand and its metal complexes
Bacterial Strains
Erwinia carotovora Proteus vulgari Bacillus stubtilis Staphylococcus aureus
No. Compound
Diameter of the Growth Inhibition Zone, mm
HL
9
7
0
0
9
0
7
8
0
0
0
8
9
7
0
0
1
2
3
4
5
6
7
[Co(L)Cl(H2O(].0.5EtOH.3.75H2O
[Cu(L)Cl(H2O)2].EtOH.2.5H2O
[Cu2(HL)3Br2(OH)(H2O)]Br.6.5H2O
[Cu(L)(OH)(H2O)].0.25EtOH.0.5H2O
[Cu (HL)(L)]ClO4.4.5EtOH
[Pd(L)Cl].5H2O
7
8
10
12
7
7
8
7
7
7
8
9
7
7
0
[Fe (HL)Cl (OH)2(H2O)].0.75EtOH.1.25H2O
Control
7
0
6
6
7
0
Streptomycin
12
12
12
absorptivity of DPPH• at (517 nm) after reaction with the
tested compound. At the lowest concentration (20 μg/mL)
the antioxidant activity of the free ligand was found to be
(48.56%) but, upon complexation, it decreased signifi-
cantly within the range (17.28% – 33%) (Table 9). It was
observed, that the free ligand has higher activity than that
of metal complexes but complex (2) has higher activity in
comparison to the other complexes.[66]
respectively. The inhibition zones were measured for the
compounds, standard drugs (Streptomycin), control
(DMSO) and tabulated in (Table 10). The results were com-
pared relative to those of the standard, hence, (Table 10)
revels that the Schiff base and its complexes show mild to
moderate activity against these microorganisms. The
results (Table 10) show that the ligand exhibits better anti-
bacterial activity against Erwinia carotovora and Bacillus
stubtilis than all investigated complexes.[10,66] The increase
in biological activity of Schiff base is due to the presence of
uncoordinated azomethine linkage (C=N).[109,110] Accord-
ing to the basis of Overton's concept and Tweedy's chela-
tion theory,[109] the results showed that the complexes
(2), (4), (5) are better antibacterial against Proteus vulgaris
and Staphylococcus aureus than the inactive ligand. Also,
the data depict that complex (3) is more active against
3.8.2 | Antibacterial activity
Schiff base ligand (HL) and its metal complexes were
tested for in vitro antibacterial against different pathogenic
gram bacteria such as Bacillus stubtilis, Staphylococcus
aureus and Erwinia carotovora, Proteus vulgaris,

EMAM ET AL.
23 of 25
Staphylococcus aureus than complexes (2), (4), (5) and (7)
relative to Sterptomycin (standard drug), while the other
complexes (1, 6) and the ligand do not exhibit activity
against this type of bacteria. So, the resulted items of
highest‐to‐lowest effective relative to Staphylococcus
aureus can be summarized as follows:
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characterization, thermal investigation and
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