Azo dye with nitrogen donor sets of atoms and its metal complexes: Synthesis, characterization, DFT, biological, anticancer and Molecular docking studies

Article abstract of DOI:10.1002/aoc.4347

An azo derivative was synthesized by coupling diazotized 2,6-diaminopyridine with p-dimethyl amino benzaldehyde and this new ligand formed a series of metal complexes with Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II) and Cd(II) salts. These complexes were characterized on the basis of elemental analyses, molar conductance, infrared spectroscopy, UV–Vis, ¹H NMR, mass spectrometry, electronic spectra, magnetic susceptibility and ESR spectral studies, conductivity measurements, thermogravimetric analyses (TG-DTG). The molecular and electronic structure of the azo ligand was optimized theoretically and the quantum chemical parameters were calculated. The ligand and its metal complexes were subjected to X-ray powder diffraction study. The thermal stability of the ligand and its metal complexes was examined by thermogravimetry. The ligand and its complexes were tested for their in vitro antimicrobial activity, some of the complexes showed good antimicrobial activities against some selected bacterial and fungal strains. Anticancer activity of the ligand and its metal complexes are evaluated against human cancer (MCF-7 cells viability). Molecular docking was used to predict the binding between azo ligand and the receptors of nucleoside diphosphate kinase of Staphylococcus aureus (3Q8U) and (3HB5) which is breast cancer mutant oxidoreductase. The docking study provided useful structural information for inhibition studies.

Full text of DOI:10.1002/aoc.4347

Received: 6 January 2018
DOI: 10.1002/aoc.4347
Revised: 7 February 2018
Accepted: 12 February 2018
F U L L P A P E R
Azo dye with nitrogen donor sets of atoms and its metal
complexes: Synthesis, characterization, DFT, biological,
anticancer and Molecular docking studies
Walaa H. Mahmoud1
| Fatma N. Sayed1 | Gehad G. Mohamed1,2
1 Chemistry Department, Faculty of
Science, Cairo University, Giza 12613,
Egypt
2 Egypt Nanotechnology Center, Cairo
University, El‐Sheikh Zayed, 6th October
An azo derivative was synthesized by coupling diazotized 2,6‐diaminopyridine
with p‐dimethyl amino benzaldehyde and this new ligand formed a series of
metal complexes with Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II) and
Cd(II) salts. These complexes were characterized on the basis of elemental anal-
12588, Egypt
1
yses, molar conductance, infrared spectroscopy, UV–Vis, H NMR, mass spec-
Correspondence
trometry, electronic spectra, magnetic susceptibility and ESR spectral studies,
conductivity measurements, thermogravimetric analyses (TG‐DTG). The
molecular and electronic structure of the azo ligand was optimized theoretically
and the quantum chemical parameters were calculated. The ligand and its
metal complexes were subjected to X‐ray powder diffraction study. The thermal
stability of the ligand and its metal complexes was examined by
thermogravimetry. The ligand and its complexes were tested for their in vitro
antimicrobial activity, some of the complexes showed good antimicrobial activ-
ities against some selected bacterial and fungal strains. Anticancer activity of
the ligand and its metal complexes are evaluated against human cancer
(MCF‐7 cells viability). Molecular docking was used to predict the binding
between azo ligand and the receptors of nucleoside diphosphate kinase of
Staphylococcus aureus (3Q8U) and (3HB5) which is breast cancer mutant oxido-
reductase. The docking study provided useful structural information for inhibi-
tion studies.
Dr. Walaa Hamed Mahmoud, Chemistry
Department, Faculty of Science, Cairo
University, Giza, 12613, Egypt.
Email: dr.walaa@yahoo.com
KEYWORDS
azo ligand, MCF‐7 cells, molecular docking, quantum chemical parameters, spectroscopy
1 | INTRODUCTION
The great majority of azo dyes are monoazo com-
pounds, which have the common structure unit of the
Azo dyes and pigments constitute by far the most
important chemical class of organic colorant owing to
their versatile application in various fields of dye and
organic chemistry.[1] Azo dyes synthesized from hetero-
cyclic compounds have attracted much attention for
their bright and strong color shades ranging from
yellow to greenish blue on synthetic and natural
fabrics.[1,2]
azo chromophore N=N linking two aromatic systems.
The textile industry is the largest consumer of dye
stuffs. Although some azo dyes have been reported to
be toxic, dozens of additional monoazo dyes are permit-
ted in drugs and cosmetics.[3] The oxidation reduction
behavior of these compounds play an important role
in its biological activity. The pharmaceutical importance
of compounds including an aryl azo group has been
Appl Organometal Chem. 2018;e4347.
wileyonlinelibrary.com/journal/aoc
Copyright © 2018 John Wiley & Sons, Ltd.
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https://doi.org/10.1002/aoc.4347

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MAHMOUD ET AL.
extensively reported in the literature.[4] The tautomer-
ism of azo benzene derivatives leads to their applicabil-
ity as biomarkers and chemical sensors.[5] Because of
the good thermal stability of azo compounds, one of
the most important applications of azo compounds is
in the optical data storage. In general, cyanine dyes,
phthalocyanine dyes, and metaleazo complex dyes are
used for DVD‐R (digital versatile disc‐recordable) as
recording layer. On the other hand, organic azo com-
pounds and metal‐azo complexes are more stable than
cyanine dyes against light, provide easier control of
the wavelength according to the substituted groups,
and have good thermal stability with a metal complex.[6]
Because of the importance of azo‐containing compounds
and in continuance of our interest in syntheses of azo‐
based compounds,[7] we report herein the syntheses
and study of new azo compounds derived from 2,6‐
diaminopyridine with p‐dimethyl amino benzaldehyde
and its transition metal complexes.
2.3 | Solutions for anticancer activity
studies
A fresh stock solution of 1x10−3 M of azodye ligand
(0.00429 g/L) was prepared in the appropriate volume of
DMF. Dimethylsulphoxide (DMSO); used in cryopreser-
vation of cells and RPMI‐1640 medium was supplied from
Sigma. The medium was used for culturing and mainte-
nance of the human tumor cell line. The medium was
supplied in a powder form. It was prepared as follows:
10.4 g medium was weighed, mixed with 2 g sodium
bicarbonate, completed to 1 l with distilled water and
shaked carefully till complete dissolution. The medium
was then sterilized by filtration in a Millipore bacterial fil-
ter (0.22 μm). The prepared medium was kept in a refrig-
erator (4 °C) and checked at regular intervals for
contamination. Before use, the medium was warmed at
37 °C in a water bath and supplemented with penicillin/
streptomycin and FBS. Sodium bicarbonate (Sigma
Chemical Co., St. Louis, MO, USA) was used for the prep-
aration of RPMI‐1640 medium. 0.05% Isotonic trypan blue
solution (Sigma Chemical Co., St. Louis, MO, USA) was
prepared in normal saline and was used for viability
counting. 10% Fetal Bovine Serum(FBS) (heat inactivated
at 56 °C for 30 min), 100 units/ml Penicillin and 2 mg/ml
Streptomycin were supplied from Sigma Chemical Co., St.
Louis, MO, USA and were used for the supplementation
of RPMI‐1640 medium prior to use. 0.025% (w/v) Trypsin
(used for the harvesting of cells), 1% (v/v) acetic acid (was
used for dissolving the unbound SRB dye), 0.4%
sulphorhodamine‐B (SRB) dissolved in 1% acetic acid
was used as a protein dye. Stock solution of trichloro
acetic acid (TCA, 50%) was prepared and stored. All of
them were supplied from Sigma. 50 μl of the stock solu-
tion was added to 200 μl RPMI‐1640 medium/well to yield
a final concentration of 10% used for protein precipitation.
100% Isopropanol and 70% ethanol were used. Tris base
10 mM (pH 10.5) was used for SRB dye solubilization.
121.1 g of tris base was dissolved in 1000 ml of distilled
water and pH was adjusted by HCl (2 M).
2 | EXPERIMENTAL
2.1 | Materials and reagents
All chemicals used were of the analytical reagent grade
(AR) and of highest purity available. The chemicals used
included 2,6‐diaminopyridine (Sigma Aldrich, Germany),
p‐dimethylaminobenzaldehyde (Sigma Aldrich, Germany),
CrCl3.6H2O (Sigma Aldrich, Germany), MnCl2.2H2O (Sigma
Aldrich, Germany), NiCl2.6H2O (BDH), FeCl3.6H2O (Sigma
Aldrich, Germany), CoCl2.6H2O (Sigma Aldrich, Germany),
CuCl2.2H2O (Merck, Germany), CdCl2 (Sigma Aldrich,
Germany) and ZnCl2 (Strem Chemicals, United States).
Organic solvents were spectroscopic pure from BDH
included ethanol and dimethyl formamide. Hydrochloric
acid, sodium nitrite and sodium acetate (A.R., Sigma
Aldrich, Germany) were used.
Human tumor cell line (Breast cell line MCF7) was
obtained frozen in liquid nitrogen (‐180 °C) from the Ameri-
can Type Culture Collection and was maintained in the
National Cancer Institute, Cairo, Egypt, by serial sub‐
culturing.
2.4 | Measurements
Microanalyses of carbon, hydrogen and nitrogen were
carried out at the Microanalytical Center, Cairo Univer-
sity, Egypt, using CHNS‐932 (LECO) Vario Elemental
Analyzer. Analyses of the metals followed the dissolution
of the solid complexes in concentrated HNO3, neutraliz-
ing the diluted aqueous solutions with ammonia and
titrating the metal solutions with EDTA. 1H NMR spectra,
as a solution in DMSO‐d6, were recorded on a 300 MHz
Varian‐Oxford Mercury at room temperature using TMS
as an internal standard. Electron spin resonance spectra
2.2 | Solutions
Stock solutions of metal complexes of 1x10−3 M were pre-
pared by dissolving the accurately weighed amount of the
complexes in DMF for measuring conductivity. Solutions
of the azo dye ligand (1x10−5 M), its Cr(III), Co(II) and
Zn(II) complexes (1x10−4 M) and its Mn(II), Fe(III),
Ni(II), Cu(II) and Cd(II) complexes (5x10−5) were pre-
pared by accurate dilution of the previous prepared stock
solutions to measure UV–Vis spectra.

MAHMOUD ET AL.
3 of 19
TABLE 1 The different optimized and quantum chemical parameters of free L ligand
Atoms
Bond lengths (Å)
Atoms
Bond lengths (Å)
C1‐C2
1.4062
1.4136
1.0837
1.4206
1.4309
1.4017
1.0843
1.4084
1.087
C25‐C27
C25‐N33
C26‐C28
C26‐H29
C27‐C30
C27‐H31
C28‐C30
C28‐C53
C30‐H32
N32‐C43
N32‐C47
N33‐35C
N33‐C39
C35‐H36
C35‐H37
C35‐H38
C39‐H40
C39‐H41
C39‐H42
C43‐H44
C43‐H45
C43‐H46
C47‐H48
C47‐H49
C47‐H50
C51‐H52
C51‐O55
C53‐H50
C53‐O54
1.4399
1.3835
1.3989
1.087
C1‐C6
C1‐H7
C2‐C3
C2‐N11
C3‐C4
1.3858
1.0834
1.4232
1.4717
1.087
C3‐H8
C4‐C5
C4‐H9
C5‐C6
1.4122
1.0862
1.4295
1.2939
1.2945
1.4203
1.4222
1.447
1.476
C5‐H10
C6‐N12
N9‐N12
N10‐N11
N11‐C15
N12‐C24
C15‐C16
C15‐C17
C16‐C18
C16‐N32
C17‐C19
C17‐H20
C18‐C21
C18‐H22
C19‐C21
C19‐C51
C21‐H23
C24,‐C25
C24‐C26
1.475
1.4753
1.4747
1.0977
1.0978
1.09
1.0941
1.1015
1.0949
1.0972
1.0974
1.0903
1.0939
1.1014
1.095
1.4154
1.4404
1.3828
1.3983
1.087
1.3854
1.0833
1.4236
1.472
1.1086
1.2536
1.1088
1.2536
1.087
1.4461
1.4147
The calculated quantum chemical parameters
E (a.u.)
−1407.61
6.73
Dipole moment (Debye)
EHOMO (eV)
ELUMO (eV)
Δ E (eV)
χ(eV)
−5.71
−2.77
2.94
4.24
η (eV)
1.47
σ (eV)−1
0.68
Pi (eV)
S (eV)−1
−4.24
0.34
ω (eV)
6.11
ΔNmax
2.88

4 of 19
MAHMOUD ET AL.
were also recorded on JES‐FE2XG ESR spectrophotome-
ter at Microanalytical Center, Tanta University. The X‐
ray powder diffraction analyses were carried out using
Philips Analytical X‐ray BV, diffractometer type PW
1840. Radiation was provided by copper target (Cu anode
2000 W) high intensity X‐ray tube operated at 40 KV and
25 mA. Divergence and the receiving slits were 1 and 0.2,
respectively. Mass spectra were recorded by the EI tech-
nique at 70 eV using MS‐5988 GS‐MS Hewlett–Packard
instrument at the Microanalytical Center, National Cen-
ter for Research, Egypt. FT‐IR spectra were recorded on
a Perkin‐Elmer 1650 spectrometer (4000–400 cm−1) in
KBr pellets. The electronic spectra were recorded in
DMSO at room temperature on Shimadzu UV–visible
mini‐1240 spectrophotometer. The molar magnetic sus-
ceptibility was measured on powdered samples using the
Faraday method. The diamagnetic corrections were made
by Pascal's constant and Hg[Co(SCN)4] was used as a
calibrant. Molar conductivities of 10−3 M solutions of
the solid complexes in DMF were measured using Jenway
4010 conductivity meter. The thermogravimetric analyses
(TG and DTG) of the solid complexes were carried out
from room temperature to 1000 °C using a Shimadzu
TG‐50H thermal analyzer. Diffused reflectance spectra
analyses were carried out at the Microanalytical Center,
Cairo University, Egypt. The antimicrobial activities were
carried out at the Microanalytical Center, Cairo Univer-
sity, Egypt. The anticancer activity was performed at the
National Cancer Institute, Cancer Biology Department,
Pharmacology Department, Cairo University. The optical
density (O.D.) of each well was measured spectrophoto-
metrically at 564 nm with an ELIZA microplate reader
(Meter tech. R 960, USA).
2.6 | Synthesis of metal complexes
Azodye ligand was dissolved in DMF (0.4 g, 9.32x10−4
mmol) and ethanolic metal salts with the same molar
ratio (0.248 g Cr(III), 0.184 g Mn(II), 0.252 g Fe(III),
0.222 g Co(II), 0.220 g Ni(II), 0.125 g Cu(II), 0.127 g Zn(II)
and 0.171 Cd(II)) were added to the azo ligand solution.
The mixture was heated under reflux for 2–3 h with stir-
ring. The precipitates were filtered off, washed with etha-
nol followed by diethyl ether and dried in a vacuum
desiccator over anhydrous CaCl2.
2.7 | Spectrophotometric studies
The absorption spectra were recorded for 1x10−5
M
solutions of the azodye free ligand, 1x10−4 M solutions
of its Cr(III), Co(II) and Zn(II) metal complexes and
5x10−5 M solutions of its Mn(II), Fe(III), Ni(II), Cu(II)
and Cd(II) metal complexes dissolved in DMF. The
spectra were scanned within the wavelength range from
200 to 700 nm.
2.8 | Pharmacology
2.8.1 | Antimicrobial activity
A filter paper disk (5 mm) was transferred into 250 ml
flasks containing 20 ml of working volume of tested
solution (100 mg/ml). All flasks were autoclaved for
20 min at 121 °C. LB agar media surfaces were inocu-
lated with four investigated bacteria (Gram positive bac-
teria: Bacillus subtilis and Staphylococcus aureus, Gram
negative bacteria: Neisseria gonorrhoeae and Escherichia
coli) and one strain of fungi (Candida albicans) by diffu-
sion agar technique, then, transferred to a saturated
disk with a tested solution in the center of Petri dish
(agar plates). All the compounds were placed at 4 equi-
distant places at a distance of 2 cm from the center in
the inoculated Petriplates. DMSO served as control.
Finally, all these Petri dishes were incubated at 25 °C
for 48 h where clear or inhibition zones were detected
around each disk. Control flask of the experiment was
designed to perform under the same condition described
previously for each microorganism but with dimethyl‐
formamide solution only and by subtracting the diame-
ter of inhibition zone resulting with dimethyl formam-
ide from that obtained in each case, so antibacterial
activity could be calculated.[8,9] Amikacin and ketocona-
zole were used as reference compounds for antibacterial
and antifungal activities, respectively. All experiments
were performed as triplicate and data plotted were the
mean value.
2.5 | Synthesis of azodye ligand
In 250 ml quickfit round bottom flask, 2,6‐
diaminopyridine (3.00 g, 27.30 mmol) was dissolved in
ethanol (50 ml). The solution was cooled with stirring in
an ice bath to 0 °C until it becomes clear, then HCl solu-
tion (4 ml) was added. Solution of sodium nitrite (3.76 g,
54.50 mmol) in distelled water (20 ml) was added dropwise
during 10 min and the reaction mixture was further stirred
for 20 min in an ice bath at 0‐5 °C. The solution was added
dropwise to p‐dimethylaminobenzaldehyde coupling com-
ponent (8.10 g, 54.50 mmol) in ethanol (30 ml) and aque-
ous solution of 4 g sodium acetate as catalyst, with
stirring in an ice bath for a further 1 h. The product was
collected by filtration and washed with water and dried
under vacuum at room temperature. The physical and
analytical data of the isolated azodye ligand are listed in
Table 3. The compound has high melting point and it is
found to be air stable.

MAHMOUD ET AL.
5 of 19
docking software[12] and is an interactive molecular
graphics program for calculating and displaying feasible
docking modes of a receptor and ligand and complex mol-
ecules. It necessitates the ligand and the receptor as input
in PDB format. The amino acid chain was kept and the
water molecules and co‐crystallized ligands were
removed. The structure of ligand in PDB file format was
created by Gaussian03 software. The crystal structures of
receptors (3Q8U) and (3HB5) were downloaded from the
protein data bank (http://www.rcsb.org./pdb).
2.9 | Anticancer activity
Potential cytotoxicity of the complexes was tested using
the method of Skehan and Storeng.[8,10] Cells were plated
in 96‐multiwell plate (104 cells/well) for 24 h before treat-
ment with the complexes to allow attachment of cell to
the wall of the plate. Different concentrations of the com-
plexes under investigation (0, 5, 12.5, 25, 50 and 100 μg/
ml) were added to the cell monolayer triplicate wells were
prepared for each individual dose. The monolayer cells
were incubated with the complexes for 48 h at 37 °C
and in 5% CO2 atmosphere. After 48 h, cells were fixed,
washed and stained with SRB stain. Excess stain was
washed with acetic acid and attached stain was recovered
with tris–EDTA buffer. The optical density (O.D.) of each
well was measured spectrophotometrically at 564 nm
with an ELIZA microplate reader and the mean back-
ground absorbance was automatically subtracted and
mean values of each drug concentration was calculated.
The relation between surviving fraction and drug concen-
tration is plotted to get the survival curve of Breast tumor
cell line for each complex.
3 | RESULTS AND DISCUSSION
3.1 | Characterization of azodye ligand
Free ligand and its metal complexes are soluble in DMF
and DMSO solvents and have high and sharp melting
points which indicate their pure nature. The infrared
spectrum of azo ligand (L) gives interesting results and
conclusions, the most effective bands are listed in
Table 4. The frequency for the N=N stretching laid in
the region 1542 cm−1 [13]
also two bands are observed
,
Calculation:
due to vibrational ν(C=N) mode of pyridine ring at
The percentage of cell survival was calculated as
follows:
1596 cm−1 and the ν(C=O) frequency in the free L at
1665 cm .
−1 [14] The IR spectrum of the ligand also showed
a sharp band at 1162 cm−1 due to vibration of υ(N(CH3)2),
also three bands were observed due to υ(C‐N) stretching
and δ(C=N) bending at 1475 cm−1and 631 cm‐1,
respectively.
Survival fraction ¼ O:D:ðtreated cellsÞ=O:D:ðcontrol cellsÞ:
The IC50 values (the concentrations of the azo ligand
or its complexes required to produce 50% inhibition of cell
growth). The experiment was repeated three times for
MCF7 cell line.
1H NMR spectrum of free ligand (L) in DMSO‐d6 show
three characteristic signals at:
3.03 ppm (s, 12H, CH3) and 9.66 ppm (s, 2H, CHO)
assigned to N(CH3)2 and aldehydic groups, respec-
tively.[15] In the aromatic region, a few doublets and in
few cases some overlapping multiplets are observed in
the range δ (6.76–7.68) ppm. These multiplets are due to
pyridine ring and aryl protons.[13] From the mass spec-
trometry data, the azodye ligand shows molecular ion
peak at m/z = 428 amu which was in close agreement
with expected molecular weight at 429 g/mol.
2.10 | Computational methodology
The electronic structure calculations of azo ligand were
carried out using the Gaussian03 suite of program.[11]
They were fully optimized employing DFT based B3LYP
method along with the LANL2DZ basis set. In order to
incorporate the effect of the solvent around the molecule,
the TD‐DFT method (along with LANL2DZ basic set) was
used to calculate the electronic absorption spectra of the
azo ligand. The contribution of molecular orbital on
HOMO and LUMO was also calculated.
3.2 | Geometry optimization
The optimized structures of the azo ligand (L) were cal-
[11]
culated using the Gaussian03 suite of program.
They
were fully optimized employing DFT based B3LYP
method along with the LANL2DZ basis set. Table 1
shows some selected geometric parameters from the
optimized structure. Figure 1 illustrates the optimized
structure of azo ligand (L). The HOMO and LUMO of
L is shown in Figure 2. The HOMO–LUMO energy
gap, ΔE, which is an important stability index, is
applied to develop theoretical models for explaining
2.11 | Molecular docking
In order to find out the possible binding modes of the
most active compounds against the receptors of nucleo-
side diphosphate kinase of Staphylococcus aureus
(3Q8U) and (3HB5) which is breast cancer mutant oxido-
reductase. The molecular docking studies were performed
using MOE 2008 software and it is rigid molecular

6 of 19
MAHMOUD ET AL.
FIGURE 1 The optimized structure of the free ligand (L)
FIGURE 2 Theoretical electronic absorption transitions for L in DMF solvent
the structure and conformation barriers in many molec-
ular systems.[16,17]
The calculated quantum chemical parameters are
given in Table 1. Additional parameters such as ΔE, abso-
lute electronegativities, χ, chemical potentials, Pi, absolute
hardness, η, absolute softness, σ, global electrophilicity, ω,
global softness, S, and additional electronic charge,
ΔNmax, have been calculated according to the following
equations:[18–20]
ΔE ¼ ELUMO−EHOMO
(1)

MAHMOUD ET AL.
7 of 19
electrical transport properties due to providing measure-
ment of electron conductivity. In addition, this energy
gap characterizes the molecular stability and spectro-
scopic properties of the molecular systems. The smaller
energy gap describes chemically soft molecule which can
be easily polarizable. The energy gap of L was calculated
at 2.94 eV by TD‐DFT level.
−ðEHOMO þ ELUMO
Þ
χ ¼
(2)
(3)
(4)
(5)
(6)
2
E
LUMO−EHOMO
η ¼
2
1
σ ¼
η
3.4 | Characterization of metal complexes
3.4.1 | Elemental analyses:
Pi ¼ − χ
The elemental analysis data of the complexes (Table 3)
show the formation of the complexes in the ratio of 1:1
for [ML]. It is found that the theoretical values are in
agreement with the found values. The metal complexes
have high melting points and they are found to be air sta-
ble. The azodye ligand is soluble in common organic sol-
vents and all the complexes are freely soluble in DMF
and DMSO but insoluble in methanol, ethanol and water.
1
S ¼
2η
Pi2
ω ¼
(7)
(8)
2η
Pi
ΔNmax ¼ −
η
3.4.2 | Molar conductance measurements:
3.3 | Electronic spectra and frontier
molecular orbital energies
The conductivity data in DMF are in the range reported
for 2:1 and 1:1 electrolytes. The complexes of [Cr(L)Cl]
Cl2.2H2O, [Mn(L)H2O]Cl2, [Fe(L)Cl]Cl2.H2O and [Cu(L)
The UV–Vis electronic spectra of L in DMF were recorded
at the range of 700–200 nm. The theoretical electronic
absorption wavelengths and oscillator strengths were cal-
culated by the TD‐DFT method with B3LYP /LanL2DZ
basis sets in above solvent. Obtained theoretical results
are summarized along with the experimental ones in
Table 2. The data shows that there are two absorption
bands at 255 and 388 nm in DMF solvent. The corre-
sponding peaks in TD‐DFT calculations are found at the
range of 253, 337, 493 and 533 nm in DMF solvent as
can be seen in Table 5. The patterns of highest occupied
(HOMO) and the lowest unoccupied molecular orbitals
(LUMO) have been given in Figure 2. While LUMO repre-
sents the ability to obtain an electron, HOMO represents
the ability to donate an electron. In other words, the
energy of the HOMO is directly related to the ionization
potential, LUMO energy is directly related to the electron
affinity. Hence these frontier energies play an important
role in the electric and optical properties as well as ener-
gies is a critical parameter in determining molecular
H2O]Cl2.H2O have values indicate that they are 2:1 elec-
[21]
trolytes and other complexes are 1:1 electrolytes
.
3.5 | IR spectral studies
IR spectral study has proven to be the most suitable tech-
nique to give enough information to elucidate the way of
bonding of ligand to the metal ions Table 4. The bands
arising due to υ(C=N) pyridine stretching at 1596 cm‐1[21]
which shifted to 1591–1602 cm−1, υ(N=N) at 1542 cm−1
was observed to be shifted to 1537–1540 cm‐1[22] and the
υ(N(CH3)2) stretching frequency in the free ligand at
1162 cm−1 that showed shift to 1160–1174 cm−1 in the
metal complexes indicating the involvement of nitrogen
of pyridine ring, nitrogen of the azo group and nitrogen
of amino group in the complex formation. The band was
observed at 1665 cm−1 assigned to υ(C=O) stretching fre-
quency showed very little changes at 1660–1664 cm−1
which indicated to non‐contribution of carbonyl group in
TABLE 2 Main calculated optical transition with composite ion in terms of molecular orbital
Compound
L
Transition
Excitation energy (ev)
λ
Calc. (nm)
λ
exp. (nm)
max
Oscillating strength
max
HOMO → LUMO
2.94
3.06
3.86
4.43
533
493
337
253
……
……
388
255
0.01
0.71
0.05
0.20
HOMO → LUMO+1
HOMO‐1 → LUMO+2
HOMO‐2 → LUMO+2

8 of 19
MAHMOUD ET AL.

MAHMOUD ET AL.
9 of 19
the coordination.[23] Also new bands were observed at
939–960 cm−1 and 815–821 cm−1 due to H2O stretch of
coordinated water in Mn(II), Cu(II) and Zn(II) complexes
and new bands at 518–522 cm−1 due to M‐O stretch of
coordinated water in the same complexes. IR spectrum
showed a new band at the region 424–495 cm−1 assigned
to υ(M‐N).[24]
3.6.3 | Electronic spectra and magnetic
moment measurements:
For the hexa‐coordinate Cr(III) complex, its diffused
reflectance spectrum exhibits three absorption bands at
21,980, 19,444 and 17,310 cm−1 which can be assigned
4
4
to the three spins allowed transitions A2g(F) → T1g(P),
4A2g(F) → 4T2 g(F) and 4A2g(F) → 4T1g(F). The electronic
spectrum of the chelate reported here is in reasonable
agreement with those in the literature. The magnetic
moment at room temperature is 3.84 B.M. which corre-
sponds to the expected value for octahedral Cr(III)
complexes.[28]
3.6 | 1H NMR spectral studies
1H NMR spectral data in DMSO solution of the ligand,
Zn(II) and Cd(II) complexes showed that the resonance
of protons has been assigned on the basis of their integra-
The diffused reflectance spectrum of Mn(II) complex
has transition bands at 26,101, 21,730 and 17,910 cm−1
which correspond to charge transfer (LMCT),
1
tion and multiplicity pattern. H NMR spectrum of the
ligand reveals the following signals at:
6
5
6A1g → T2g(G) and A1g → T1g, respectively, indicating
the octahedral geometry. The magnetic moment value of
the Mn(II) complex is 5.64 B M, normally observed for
octahedral coordination geometry.[29]
δ = 9.66 ppm (s, 2H, CHO), 6.76–7.68 ppm (m, 9H,
pyridine ring and two benzene rings) and 3.03 ppm (s,
12H, CH3).[25]
While, the 1H NMR spectrum of Zn(II) reveals the fol-
lowing signals at δ = 9.64 (s, 2H, CHO), 6.75–7.67 (m, 9H,
pyridine ring) and 3.02 (s, 12H, CH3) and spectrum of
Cd(II) complex showed signals at δ = 9.64, 9.63 (s, 2H,
CHO), 6.75–7.67 (m, 9H, pyridine ring) and 3.02, 3.03 (s,
12H, CH3). This shifting indicated that metal complexes
were formed by binding of metal ion with N(CH3)2, nitro-
gen of pyridine ring and azo group.
The diffused reflectance spectrum of Fe(III) complex
exhibits three bands at 21,740, 19,860 and 16,120 cm−1
which may be assigned to 4T2g(G) → 6A1g, 4T2g(G) → 6A1g
and 4T1g(D) → 6A1g transitions, respectively. The magnetic
moment value is found to be 5.32 B.M. which indicates the
presence of Fe(III) complex in octahedral geometry.[30,31]
The Co(II) chelate showed two bands at 22,211 and
4
4
19,435 cm−1 which can be assigned to T1g(F) → T2g (P)
4
4
and T1g(F) → A2g(F) transitions, respectively, which
suggests an octahedral geometry.[26] The magnetic suscep-
tibility measurement lies at 5.10 B.M. which is an indica-
tive of octahedral geometry.
3.6.1 | Mass spectral studies:
The electronic mass spectrum of Fe(III) and Ni(II) com-
plexes confirm formation of these complexes by showing
molecular ion peaks at m/z = 606 and 575, respectively,
which are very close to expected molecular weight which
were 610 g/mol for Fe(III) and 577 g/mol for Ni(II) and
also show peaks at m/z = 430, 432 amu, respectively, indi-
cates to the ligand. The intensities of these peaks show the
stabilities of the complexes.
The diffused reflectance spectrum of the Ni(II) complex
exhibits transition bands at 27,125, 20,185 and 12,587 cm−1
indicating that the Ni(II) complex is octahedral. The room
temperature magnetic moment of Ni(II) complex (3.147 B.
M.), corresponding to two unpaired electrons and in the
range of octahedral Ni(II) complexes.[30–33]
The magnetic moment value for the Cu(II) complex
recorded at room temperature is 1.75 B.M. (Table 1) and
corresponding to one unpaired electron. The electronic
spectrum of the Cu(II) complex gave a band at
17,985 cm−1, suggesting the existence of a transition from
dxy, dz2 and dxz, dyz transfer to the antibonding and half‐
filled dx2‐y2 level which is consistent with an octahedral
configuration.[34] The Zn(II) and Cd(II) complexes are
diamagnetic. According to the empirical formulae, an octa-
hedral geometry was suggested for these metal chelates.
3.6.2 | UV–Vis absorption studies:
UV–visible spectral data of the free ligand shows absorp-
tion bands at 338 and 255 nm corresponding to n → π*
transitions and π → π* transitions of C=O, heterocyclic
moieties and benzene ring, respectively.
In the spectra of metal ions complexes these bands are
shifted to higher or lower frequencies where the n → π*
transition bands are shifted to 397–308 nm in all com-
plexes and π → π* transition bands are shifted to 268–
251 nm in Mn(II), Co(II), Ni(II), Zn(II) and Cd(II), while
disappeared in Cr(III), Fe(III) and Cu(II) indicating that
the ligands coordinate to metal ions.[26,27]
3.6.4 | ESR studies:
The ESR spectrum of the Cu(II) complex was recorded in
DMSO at 300 and 77 K and its parameters and showed

10 of 19
MAHMOUD ET AL.
that the observed order for Cu(II) complex (A|| (197) > A
(98.88); g|| (2.04) > g (1.87) indicates that the complex
exerts an octahedral geometry.[19] The observed value of
G for the Cu(II) complex (G = 1.337) implies that the
exchange coupling is not present and misalignment is
appreciable. The trend g|| > g > ge (2.0023) shows that
the unpaired electron is localized in the dx2‐ y2 orbital of
the Cu(II) ion in complex.[17,31,35] The giso (2.04) value less
than 2.3 indicates the covalent character of the metal
ligand bond and the α2 value (1.87) suggestive of in‐plane
covalency. The calculated value (g||/A||) 124 cm−1 for the
complex is consistent with slightly distorted structure
and showed poor in‐plane π bonding.[36–39]
1
┴
δ ¼
:
ξ2
┴
The values of ξ (2.92–2.93 mm) and δ (0.116–
0.117 mm−2) indicated that the ligand and Cr(III) and
Ni(II) complexes have cryslline nature while other com-
plexes showed amorphous nature.
┴
3.7.1 | Thermal analysis studies:
The thermal stability of the complexes was investigated
using TGA (Figure 3). The thermogravimetric analysis
(TGA) data (Table 5) were obtained at a heating rate of
10 °C/min in a nitrogen atmosphere over a temperature
range of 20–1000 °C. The free ligand decomposed
completely leaving no residues in one step at the temper-
ature range 70‐400 °C. Decomposition of Cr(III) complex
occurred in five steps. The first step occurred in the tem-
perature range 35‐150 °C with maximum temperature
51 °C due to loss of H2O and 3/2Cl2 molecules, the second
step occured within the temperature range of 150‐320 °C
with Ts 184 °C corresponding to loss of C16H14N4 and
the last three steps occurred within temperature range
320‐980o C with maximum temperatures at 351, 636 and
3.7 | Powder X‐ray diffraction
spectroscopy
The average crystallite size (ξ) can be calculated from the
XRD pattern according to Debye–Scherrer equation
Kλ
1=2 cosθ
ξ ¼
β
The equation uses the reference peak width at angle
(θ), where is wavelength of X‐ray radiation
λ
1
1
(1.541874 Å), K is constant taken as 0.95 for organic com-
pounds and β1/2 is the width at half maximum of the ref-
erence diffraction peak measured in radians. The
dislocation density,δ, is the number of dislocation lines
per unit area of the crystal. The value of δ is related to
the average particle diameter (ξ) by the relation:[8,9,40]
898 °C due to loss of C7H9N3 and /4O2 leaving /2Cr2O3
as a residue. The TG curve of [Mn(L)H2O]Cl2 showed
decomposition through six steps, the first and second
steps ocurred within temperature range 35–190 °C with
maximum temperatures 152 °C and 187 °C due to loss
of H2O and ½Cl2 molecules. The third step ocurred
FIGURE 3 Thermal analyses (TG and DTG) of (a) L, (b) Cr(III), (c) Mn(II), (d) Fe(III), (e) Co(II), (f) Ni(II), (g) Cu(II), (h) Zn(II) and (i)
Cd(II) complexes

MAHMOUD ET AL.
11 of 19
TABLE 5 Thermoanalytical results (TG and DTG) of L and its metal complexes
TG range
(°C)
DTGmax
(°C)
Mass loss Total mass
loss Estim (Calcd) %
Complex
n*
Assignment
Residues
L
70–400
221
1
99.9 (100)
‐Loss of C23H23N7O2.
‐‐‐‐‐‐‐‐
[Cr(L)Cl]Cl2.H2O
35–150
150–320
320–980
51
184
351, 636, 898
1
1
3
19.5 (20.5)
44.5 (43.3)
25.3 (23.6) 87.3 (87.1)
‐Loss of H2O and 3/2Cl2.
‐Loss of C16H14N4.
‐Loss of C7H9N3 and /4O2.
1/2Cr2O3
1
[Mn(L)H2O]Cl2
[Fe(L)Cl]Cl2.H2O
[Co(L)Cl]Cl.H2O
[Ni(L)Cl]Cl.H2O
[Cu(L)H2O]Cl2.H2O
[Zn(L)Cl]Cl.H2O
[Cd(L)Cl]Cl
35–190
190–260
260–1000
152, 187
218
265, 719, 851
2
1
3
11.6 (12.1)
19.9 (18.2)
45.4 (46.8) 77.2 (77.3)
‐Loss of H2O, ½Cl2 and CH4.
‐Loss of C4H12N2O.
‐Loss of C13H7N5Cl.
MnO + 5C
30–197
198–485
486–890
64
220, 394
555, 813
1
2
2
3.6 (3.0)
21.2 (22.0)
53.8 (53.3) 78.8 (80.97)
.‐Loss of H2O.
3
‐Loss of /2Cl2 and CO.
½Fe2O3 + 3C
CoO + 3C
‐Loss of C16H23N7.
30–210
210–350
350–950
85
294
357, 685, 894
1
1
3
9.9 (9.3)
8.9 (8.8)
62.8 (62.7) 82.0 (80.8)
‐Loss of H2O and ½Cl2.
‐Loss of ½Cl2 and CH3.
‐Loss of C19H20N7O2.
30–165
165–255
255–900
52, 90
184
261, 549, 834
2
1
3
30.1 (31.0)
33.7 (33.1)
22.4 (21.5) 80.7 (80.8)
‐Loss of Cl2, H2O and C4H14N2.
‐ Loss of C12H5N3.
‐Loss of C4H4N4O.
NiO + 3C
CuO + 4C
ZnO
40–205
205–355
355–1000
80, 129
240
362, 723
2
1
2
5.9 (6.0)
10.7 (11.8)
62.0 (61.0) 79.0 (78.7)
‐Loss of 2H2O.
‐Loss of Cl2.
‐Loss of C19H23N7O.
35–130
130–340
340–770
62
321
380, 519, 683
1
1
3
11.6 (11.9)
9.5 (8.8)
64.4 (63.1) 86.0 (86.1)
‐Loss of H2O, ½Cl2 and CH4.
‐Loss of ½Cl2 and CH4.
‐Loss of C21H15N7O.
30–270
270–630
202
316, 568, 602
1
3
69.4 (69.2)
8.2 (7.6) 79.6 (79.7)
‐Loss of Cl2 and C21H15N7.
‐Loss of C2H8O.
CdO
n* = number of decomposition steps
within temperature range 190–260 °C with Ts 218 °C with
mass loss corresponding to and C4H12N2O. The last three
steps occurred within temperature range 260‐1000 °C
with maximum temperatures 265, 719, 851 °C due to loss
of C13H7N5Cl leaving MnO contaminated with carbon
atoms as residues.
range 350‐950 °C with maxima 357, 685 and 894 °C leav-
ing contaminated cobalt oxide.
The [Ni(L)Cl]Cl.H2O complex is thermally
decomposed through six successive decomposition steps.
The first two steps occurred in the range 30‐165 °C with
maxima temperatures 52 and 90 °C may be attributed to
the loss of Cl2, H2O and C4H14N2. The third step occurred
at 165‐255 °C with Ts 184 °C due to loss of C12H5N3.
Finally the last three steps of decompositions occurred
within temperature range 255‐900 °C with maxima 261,
549 and 834 °C due to loss of C4H4N4O leaving NiO + 3C.
For [Cu(L)H2O]Cl2.H2O complex the TG curve
showed mass loss in the temperature range of 40–205 °C
associated with one DTG peak at 129 °C. This step can
be assigned to loss of two water molecules in one step.
The second step occurred within range 205–355 °C with
DTGmax 240 °C due to mass loss corresponded to chlorine
gas. On further heating, the DTG registered peaks at
362 °C and 723 °C within the temperature range of 355–
1000 °C that may correspond to loss of remaining ligand
fragment (C19H23N7O) leaving CuO contaminated with
carbon atoms as a residue. The TG curve of [Zn(L)Cl]Cl.
The thermal decomposition of [Fe(L)Cl]Cl2.H2O com-
plex proceeded via five degradation steps. The first step
occurred within temperature range 30‐197 °C with maxi-
mum temperature 64 °C due to loss of water molecule,
the second and third steps of decomposition occurred
within temperature range 198‐486 °C at Ts 220 and
3
394 °C due to loss of /2 Cl2 and CO gases and finally
the last two steps occurred within temperature range
487‐1000 °C corresponding to loss of C16H23N7 leaving
½Fe2O3 + 4C residues. The TG curve of Co(II) complex
showed five decomposition steps of the complex. The first
step of decomposition occurred in the range of tempera-
ture 30‐210 °C at Ts 85 °C due to loss of water molecule
and chloride atom. The second step occurred within tem-
perature range from 210 °C to 350 °C with maximum tem-
perature 294 °C. The last three steps occurred within the

12 of 19
MAHMOUD ET AL.
H2O complex showed thermal decomposition peak within
temperature range 35–130 °C with Ts 62 ° C due to one
step of decomposition corresponded to loss of H2O,
½Cl2 and CH4 molecules. The second step occurred
within range of 130–340 °C with maximum temperature
321 °C due to loss of ½Cl2 and CH4 molecules. The last
three steps occurred at 340‐770 °C with maxima 380, 519
and 683 °C due to loss of the remaining complex leaving
ZnO a residue.
The [Cd(L)Cl]Cl complex is thermally decomposed in
four successive decomposition steps. The first step
occurred within 30‐270 °C with DTGmax 202 °C, mass loss
corresponding to Cl2 and C21H15N7. The other three steps
occurred in the range from 270 °C to 630 °C with maxima
316, 568 and 602 °C due to loss of the remaining ligand
leaving CdO residue.
Coats–Redfern[35] and Horowitz–Metzger models equa-
tions.[41] The data are summarized in Table 6.
(a) Coats‐Redfern equation
The Coats‐Redfern equation which is a typical integral
method, can be represented as:
ꢀ
n
α
T2
∫
dα=ð1−αÞ ¼ ðA=αÞ∫T1 expð−E*=RTÞdt
0
For convenience of integration, the lower limit T1 is
usually taken as zero. This equation on integration gives:
For convenience of integration, the lower limit T1 is usu-
ally taken as zero. This equation on integration gives:
ꢁ
ꢂ
ꢃ
ꢄ
ꢁ
ꢂ
ln − lnð1−αÞ=T2 ¼ −E*=RT þ ln AR=φE*
A plot of left‐hand side (LHS) against 1/T was drawn,
E* is the energy of activation in kJ mol−1 and calculated
from the slope and A in (s−1) from the intercept. The
entropy of activation ΔS* in (J K−1 mol−1) was calculated
by using equation:
3.8 | Calculation of activation
thermodynamic parameters
The kinetic parameters such as enthalpy of activation
(ΔH*), activation energy (ΔE*), free energy change of
the decomposition (ΔG*), and entropy of activation
(ΔS*) were evaluated graphically by employing the
ΔS* ¼ R ln ðAh=KBTsÞ
TABLE 6 Thermodynamic data of the thermal decomposition of L and its metal complexes
Complex
Decomp.Temp. (оC) E* (kJmol−1
)
A (s−1
)
ΔS* (KJmol−1
)
ΔH* (kJmol−1
)
ΔG* (kJmol−1
)
L
70–400
110.9 (124.3)
2.66x1011 (1.64x1011) −18.60 (−22.80) 110.5 (123.9)
183.0 (191.0)
[Cr(L)Cl]Cl2.2H2O
35–150
150–320
18.29 (18.39)
21.90 (21.71)
8.75x109 (6.87x109) −52.90 (−54.60) 68.27 (68.36)
5.25x108 (1.96x108) −44.10 (−45.20) 91.86 (91.67)
94.0 (96.80)
137.0 (140.0)
[Mn(L)H2O]Cl2
[Fe(L)Cl]Cl2.H2O
[Co(L)Cl]Cl.H2O
[Ni(L)Cl]Cl.H2O
35–130
130–260
260–1000
5.67 (4.86)
14.31 (16.06)
11.73 (11.19)
3.58x107 (1.69x107) −12.00 (−11.20) 35.64 (34.83)
4.32x105 (6.35x105) −25.40 (−30.40) 142.7 (160.2)
4.46x1010 (2.31x1010) −17.90 (−18.40) 116.8 (111.5)
3.87x108 (2.79x108) −14.70 (−14.90) 58.40 (57.60)
1.32x106 (8.36x106) −28.90 (−32.50) 121.8 (123.2)
6.32x109 (1.76x109) −46.60 (−53.60) 144.7 (146.2)
63.42 (63.81)
180.0 (180.9)
204.0 (202.3)
30–335
335–485
485–890
5.86 (5.78)
21.84 (23.26)
41.79 (46.34)
77.21 (75.68)
151.0 (152.0)
186.1 (188.0)
30–210
210–350
350–950
12.22 (11.43)
23.74 (24.25)
26.84 (26.75)
9.27x107 (7.74x107) −30.50 (−29.70) 72.19 (74.00)
1.80x1011 (2.49x1011) −37.90 (−40.50) 93.69 (94.21)
8.10x105 (5.11x105) −38.80 (−40.10) 99.79 (98.67)
129. (127.8)
121.8 (120.6)
133.0 (134.2)
30–165
165–255
255–900
13.19 (13.24)
14.85 (14.51)
27.47 (28.77)
2.51x109 (4.92x109) −32.60 (−34.20) 73.16 (73.22)
5.64x106 (6.97x106) −28.70 (−29.52) 64.81 (64.47)
1.31x106 (3.57x108) −47.50 (−51.60) 97.43 (98.73)
130.0 (129.0)
122.3 (101.0)
207.8 (207.9)
[Cu(L)H2O]Cl2.H2O 40–205
205–350
15.68 (16.31)
24.10 (23.83)
33.34 (32.97)
7.50x109 (9.47x109) −35.20 (−38.30) 75.64 (76.27)
7.28x108 (5.37x108) −43.20 (−44.20) 84.06 (83.79)
4.01x107 (1.54x107) −48.50 (−49.60) 93.29 (93.86)
149.0 (147.9)
128.0 (130.2)
124.7 (124.2)
350–1000
[Zn(L)Cl]Cl. H2O
[Cd(L)Cl]Cl
35–130
130–340
340–770
5.29 (5.37)
22.26 (23.10)
29.55 (30.62)
3.19x108 (3.35x108) −12.20 (−13.70) 42.60 (43.40)
8.46x1010 (1.20x1010) −33.60 (−36.60) 72.21 (73.50)
1.78x107 (7.17x107) −41.40 (−44.50) 129.5 (130.6)
88.23 (89.75)
122.2 (123.3)
147.2 (147.9)
30–270
270–630
10.82 (11.10)
27.11 (23.31)
6.52x1011 (7.40x1011) −18.90 (−21.00) 78.12 (76.66)
6.50x107 (4.83x107) −42.10 (−37.93) 107.6 (108.7)
118.0 (107.0)
132.8 (132.0)
#The data between parenthesis are obtained using Horowitz‐Metzger equation while the data without parenthesis are obtained using Coats‐Redfern equation.

MAHMOUD ET AL.
13 of 19
Where KB is the Boltzmann constant, h is the Plank's
constant and Ts is the DTG peak temperature.
(b) Horowitz‐Metzger equation
The Horowitz‐Metzger equation is an illustrative of the
approximation methods. These authors derived the relation:
ꢁꢃ
ꢄ
ꢂ
ꢃ
ꢄ
1−n
=ð1−nÞ ¼ E*θ=2:303RTg for n≠1
2
log 1−ð1−αÞ
When n = 1, the LHS of equation 4 would be log[−log
(1‐α)]. For a first‐order kinetic process the Horowitz‐
Metzger equation may be written in the form:
ꢁ
ꢃ
ꢄꢂ
ꢃ
ꢄ
log log Wα=Wγ ¼ E*θ=2:303RTs – log 2:303
2
Where θ = T‐Ts, wγ = wα –w, wα = mass loss at the
completion of the reaction; w = mass loss up to time t.
The plot of log[log(wα /wγ)] versus θ was drawn and found
to be linear from the slope of which E* was calculated. The
pre‐exponential factor, A, was calculated from the equation:
SCHEME 1 The proposed structures of different metal complexes
1
h
i
spectroscopy, UV–Vis, H NMR and mass spectrometry
ꢃ
ꢄ
E*=RTs ¼ A= φ expð−E*=RTsÞ
2
all confirm complexes formation. On the basis of elec-
tronic spectral data and magnetic susceptibility measure-
ments, octahedral geometry was proposed for metal
complexes. X‐ray diffraction study showed that the ligand,
Cr(III) and Ni(II) complexes have crystalline, while other
complexes have amorphous structures. The thermal sta-
bility of the ligand and its metal complexes was examined
by thermogravimetry.
The entropy of activation, ΔS*, enthalpy of activation,
ΔH* and Gibbs free energy, ΔG*, were calculated using
the following equations;
ΔH* ¼ E*–RT
ΔG* ¼ ΔH*–TΔS*
3.10 | Biological activities
The activation energies of decomposition appeared in
the range of 4.86–124.30 kJ mol−1 [25,42]
The positive sign
.
The biological activities of the free ligand and its metal
complexes were examined. A comparative study of the
growth inhibition zone values of ligand and its metal
complexes indicated that some metal complexes exhibited
higher activity than the free ligand Table 7, Figure 4.
For Bacillus subtilis (gram positive bacteria):
of ΔG* for the investigated complexes revealed that the
free energy of the final residue was higher than that of
the initial compound, and all the decomposition steps
were non‐spontaneous processes. The values of the activa-
tion, ΔG* increased significantly for the subsequent
decomposition stages of a given complex. This can be
explained based on increasing the values of TΔS* signifi-
cantly from one step to another overrides the values of
ΔH*. On the other hand, the negative values of ΔS* for
the degradation process showed more ordered activated
complex than the reactants or the reaction was slow.[43,44]
[Zn(L)H2O]Cl2
>
[Co(L)Cl]Cl.H2O
>
[Cd(L)Cl]
Cl > [Ni(L)Cl]Cl.H2O = [Cu(L)H2O]Cl2.H2O > [Mn(L)
H2O]Cl2 > [Fe(L)Cl]Cl2 = L > Amikacin > [Cr(L)Cl]Cl2.
For Staphylococcus aureus (gram positive bacteria):
[Zn(L)H2O]Cl2
=
[Cd(L)Cl]Cl
>
>
[Co(L)Cl]Cl.
[Ni(L)Cl]Cl.
H2O [Cu(L)H2O]Cl2.H2O
>
H2O > L = [Cr(L)Cl]Cl2 = [Mn(L)H2O]Cl2 = [Fe(L)Cl]
Cl2 = Amikacin.
For Escherichia coli (gram negative bacteria): [Co(L)
Cl]Cl.H2O = [Cd(L)Cl]Cl > [Zn(L)H2O]Cl2 > [Ni(L)Cl]
3.9 | Structural interpretation
The general formula of metal ions complexes can be
expected and illustrated in scheme 1, which were also
confirmed by elemental analyses that show formation of
complex in 1:1 ratio, molar conductance measurements
show that all complexes are electrolytes, infrared
Cl.H2O
> [Cu(L)H2O]Cl2.H2O > L = [Fe(L)Cl]
Cl2 > [Mn(L)H2O]Cl2 > Amikacin > [Cr(L)Cl]Cl2.
For Neisseria gonorrhoeae (gram negative bacteria):
[Zn(L)H2O]Cl2
>
[Cd(L)Cl]Cl
>
[Cu(L)H2O]Cl2.

14 of 19
MAHMOUD ET AL.
TABLE 7 Biological activity of L and its metal complexes
Inhibition zone diameter (mm / mg sample)
Gram positive Gram negative
Bacillus subtilis Staphylococcus aureus Escherichia coli Neisseria gonorrhoeae Candida albicans
fungus
Sample
Control: DMSO
L
0
9
0
9
0
10
0
0
9
0
9
[Cr(L)Cl]Cl2.2H2O
[Mn(L)H2O]Cl2
[Fe(L)Cl]Cl2.H2O
[Co(L)Cl]Cl.H2O
[Ni(L)Cl]Cl.H2O
0
9
0
0
10
9
9
9
0
0
9
10
17
13
12
15
17
7
10
14
13
16
26
20
6
0
17
11
21
10
13
28
28
9
12
14
9
[Cu(L)H2O]Cl2.H2O 11
[Zn(L)Cl]Cl.H2O
[Cd(L)Cl]Cl
Amikacin
20
16
6
0
13
‐
Ketokonazole
‐
‐
‐
‐
9
FIGURE 4 (a) Biological activity of azo
dye ligand (L) and its complexes. (b)
Relationship between inhibition zone
diameter and atomic mass of metals
H2O > [Co(L)Cl]Cl.H2O > [Ni(L)Cl]Cl.H2O > [Fe(L)Cl]
Cl2 > L > Amikacin > [Cr(L)Cl]Cl2 = [Mn(L)H2O]Cl2.
For Candida albicans (fungus): [Ni(L)Cl]Cl.
H2O > [Cd(L)Cl]Cl > [Co(L)Cl]Cl.H2O > L = [Cu(L)
H2O]Cl2.H2O = Ketokonazole > [Cr(L)Cl]Cl2 = [Mn(L)
H2O]Cl2 = [Fe(L)Cl]Cl2 = [Zn(L)H2O]Cl2.
Higher activity of metal complex was probably due
greater lipophilic nature of the complex. It increased activ-
ity of the metal complex, and can be explained on the basis
of Overtone's concept and Tweedy's chelation theory.[45]
According to Overtone's concept of cell permeability, the
lipid membrane that surrounds the cell favours the

MAHMOUD ET AL.
15 of 19
passage of only lipid soluble materials due to which
liposolubility was considered to be an important factor
that controls the antimicrobial activity. On chelation, the
polarity of the metal ion will be reduced to a greater extent
due to the overlap of the ligand orbital and partial sharing
of positive charge of metal ion with donor groups. Further,
it increases the delocalization of the π electrons over the
whole chelate ring and enhances the lipophilicity of the
complex. This increased lipophilicity enhances the
penetration of the complexes into lipid membrane and
thus blocks the metal binding sites on enzymes of microor-
ganisms. These metal complexes also disturb the respira-
tion process of the cell and thus block the synthesis of
proteins, which restricts further growth of the organ-
ism.[10,13,45–49]
The activity indexes of the newly synthesized ligand
and its metal complexes were calculated and plotted in
figure (5) according to the following equation:
FIGURE 5 Activity index of azo dye ligand and its complexes against microbial organisms
TABLE 8 Antibreastic cancer activity of L and its metal complexes
Surviving fraction (MCF7)
Complex
Concn. (μg/ ml)
0.0
5.0
12.5
25.0
50.0
IC50 (μg/ ml)
L
1.00
1.00
1.00
1.00
1.00
1.00
0.97
1.00
0.80
0.91
0.85
0.81
0.71
0.82
0.72
0.76
0.62
0.54
0.76
0.61
0.53
0.76
0.47
0.47
0.43
0.45
0.29
0.48
0.46
0.32
44.50
42.70
28.00
47.90
22.60
19.40
[Cr(L)Cl]Cl2.2H2O
[Co(L)Cl]Cl.H2O
[Ni(L)Cl]Cl.H2O
[Zn(L)Cl]Cl.H2O
[Cd(L)Cl]Cl

16 of 19
MAHMOUD ET AL.
changes in the chelation molecular structure could lead
to profound differences in anticancer activity.[50–52]
Inhibition zone of compoundðmmÞ
Activity indexðAÞ ¼
x 100
Inhibition zone of standard drugðmmÞ
The figures showed that Zn(II) complex has the
highest activity index, while Cr(III) has no activity index
against Bacillus subtilis, Escherichia coli and Neisseria
gonorrhoeae.
3.11 | Binding to biological
macromolecules
DNA is considered a major biological target for com-
pounds that have pharmacological activity. These com-
pounds are a well‐known for cancer chemotherapy.[53]
We have computationally investigated the reaction
between the azo ligand (L) and two macromolecules
(MMs). The calculated structure of L was docked with
two possible biological targets: the receptors of nucleoside
diphosphate kinase of Staphylococcus aureus (PDB code:
3Q8U), and breast cancer mutant oxidoreductase (PDB
code: 3HB5). The structures of these MMs were obtained
from the Protein Data Bank. Table 9 contains the calcu-
lated binding energies for the interaction with these
MMs. Generally, the azo ligand demonstrated strong
binding interactions with two investigated proteins as
demonstrated by the significantly negative binding ener-
gies computed in each case. It was observed that the azo
ligand bound stronger to 3Q8U than to 3HB5 with a low
binding energy −3.3 kcal/mol. Representative binding
3.10.1 | Anticancer activity evaluation:
The MCF7 cell line was chosen to study antibreastic can-
cer effect of the newly synthesized ligand and its metal ion
complexes. The values of concentration at which half of
the maximal effect is observed (IC50) with the numerical
data summarized in Table 8. Figure 6 showed relation
between surviving fraction and concentration of com-
pounds under investigation. It was clear from Table 8 that
the free ligand has IC50 value of 44.5 μg/ml, while Cr(III),
Co(II), Ni(II), Zn(II) and Cd(II) complexes have IC50 in
the range of 19.4–47.9 μg/ml. the highest IC50 value was
of Ni(II) complex (47.9 μg/ml), the lowest value was of
Cd(II) complex (19.4 μg/ml), this means that Cd(II) com-
plex had very high cytotoxic activity against MCF‐7 cell
line. These results provided a perfect example of how
FIGURE 6 Anticancer activity of azo dye ligand and its complexes
TABLE 9 Energy values obtained in docking calculations of L with receptors of crystal structure of breast cancer mutant 3HB5‐oxidore-
ductase and crystal structure of S. aureus (3Q8U)
Receptor
Ligand moeity
Receptor site
Interaction
Distance (Aο)
E (kcal/mol)
3HB5
O 56
6‐ring
6‐ring
N MET 193
CD1 ILE 14
CD2 PHE 192
H‐acceptor
pi‐H
pi‐H
2.92
3.78
3.65
−0.8
−0.8
−0.7
3Q8U
N 12
O 55
6‐ring
6‐ring
N THR 91
H‐acceptor
H‐acceptor
pi‐cation
pi‐H
3.30
3.18
3.96
4.38
−0.1
−3.3
−0.7
−1.1
NH1 ARG 125
NH1 ARG 85
N THR 91

MAHMOUD ET AL.
17 of 19
FIGURE 7 The interaction between i) [L and receptor 3Q8U (a) 2D plot, (b) 3D plot] and ii)[L and receptor 3HB5 (c) 2D plot and (d) 3D
plot]
poses for the azo ligand revealed the strongly favorable
affinities to be a consequence of extensive hydrophobic
interactions with binding site residues Figure 7.
temperature showed axial type (dx2‐y2) with covalent
bond character in an octahedral environment.
The proposed metal ions complexes structures as
follow:
½MLClꢀClx:yH2O where M ¼ CrðIIIÞ; FeðIIIÞ;
x ¼ 2; y ¼ 0; M ¼ CoðIIÞ; NiðIIÞ;
x ¼ 1; y ¼ 1 and M ¼ CdðIIÞ;
4 | CONCLUSION
In this work, a new free azodye ligand was prepared by
coupling diazotized 2,6‐diaminopyridine with p‐
dimethylaminobenzaldehyde then a series of metal ions
complexes were synthesized. The free ligand and its com-
plexes were characterized using different techniques. Ele-
mental analysis showed that formation of ligand:metal in
ratio 1:1. IR analysis confirmed coordination of ligand to
metal via pyridine nitrogen, azo nitrogen and tertiary
x ¼ 1; y ¼ 0:½MLðH2OÞꢀClx:yH2O
where M ¼ MnðIIÞ; ZnðIIÞ; x ¼ 2; y ¼ 0 and
M ¼ CuðIIÞ; x ¼ 2; y ¼ 1:
1
amine nitrogen also HNMR confirmed that. Mass spec-
trometry showed molecular ion peak as the expected
molecular weight. Conductivity measurements showed
that all metal complexes are electrolytes. The magnetic
and solid reflectance measurements confirmed octahedral
geometry of the complexes, and XRD results showed that
the free ligand, Cr(III) and Ni(II) complexes have crystal-
line structure while other complexes have amorphous
nature. ESR spectra of solid Cu(II) complex at room
The antimicrobial studies showed that most metal
ions complexes have higher activity than the free ligand.
The antibreastic cancer test showed that Ligand and
complexes have high IC50 but Ni(II) complex showed
the highest IC50 (47.90 μg/ ml). The docking study with
two different proteins provided useful structural informa-
tion for inhibition studies of free azo ligand.

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MAHMOUD ET AL.
[18] W. H. Mahmoud, N. F. Mahmoud, G. G. Mohamed, A. A. El‐
Bindary, A. Z. El‐Sonbati, J. Mol. Struct. 2012, 1086, 266.
ORCID
Walaa H. Mahmoud
4325
http://orcid.org/0000-0001-9187-
[19] A. Z. El‐Sonbati, M. A. Diab, A. A. El‐Bindary, A. M. Eldesoky,
S. M. Morgan, Spectrochim. Acta A 2015, 135, 774.
[20] N. A. El‐Ghamaz, M. A. Diab, A. A. El‐Bindary, A. Z. El‐
Sonbati, H. A. Seyam, J. Mater. Sci. Semicond. Proc. 2014,
27, 521.
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FN, Mohamed GG. Azo dye with nitrogen donor
sets of atoms and its metal complexes: Synthesis,
characterization, DFT, biological, anticancer and
Molecular docking studies. Appl Organometal
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