Synthesis, spectral, DFT calculations and biological studies of solvatochromic copper(II)-ONS hydrazone derived from 2-aminochromone-3-carboxaldehyde

Article abstract of DOI:10.1002/aoc.5763

New ONS hydrazone ligand, 2-[(2-aminochromon-3-yl)methylidene]-N-phenylhydrazinecarbothioamide, HL, was synthesized and reacted with different salts of Cu (II) ion (OAc^?, NO₃^?, SO₄^2? and Cl^?) in absence and presence of secondary ligands (L′); 8-hydroxyquinoline, 1,10-phenanthroline or SCN^?; to form binary and ternary Cu(II)-chelates. The ligand and its Cu(II)-complexes were fully characterized by analytical, spectral, thermal, conductivity and magnetic susceptibility measurements. The metal chelates showed octahedral, square planar and /or distorted tetraherdal arrangements. Coats–Redfern equations used to calculate the kinetic parameters of the thermal decomposition stages (E_a, A, ΔH, ΔS and ΔG). The compounds exhibit luminescence property; promising interesting potential applications as photoactive materials. Lippert–Mataga, Bakhshiev, Kawski–Chamma–Viallet and microscopic solvent polarity parameter and E_T^N correlation methods were applied on the solvatochromic shifts of emission spectra to evaluate the ground (μ_g) and excited (μ_e) states dipole moments. Excited state dipole moment is larger than the ground state which may be attributed to π-π* transition. The coordinating anions play an important role on the position and intensity of emission band. The ligand and its metal complexes showed antimicrobial activity towards Gram–positive bacteria, Gram–negative bacteria, yeast and fungus. The molecular structural parameters of HL and its Cu(II)- complexes have been calculated on the basis of DFT engaged in the Gaussian 09 program at the B3LYP/6-31G(d,p) level; the theoretical data are correlated with the experimental data.

Full text of DOI:10.1002/aoc.5763

Received: 2 March 2020
DOI: 10.1002/aoc.5763
Revised: 9 April 2020
Accepted: 10 April 2020
F U L L P A P E R
Synthesis, spectral, DFT calculations and biological studies
of solvatochromic copper(II)-ONS hydrazone derived from
2-aminochromone-3-carboxaldehyde
Omima M.I. Adly
| Ali Taha | Magdy A. Ibrahim | Shery A. Fahmy
Department of Chemistry, Faculty of
Education, Ain Shams University, Roxy,
Cairo, Egypt
New ONS hydrazone ligand, 2-[(2-aminochromon-3-yl)methylidene]-N-
phenylhydrazinecarbothioamide, HL, was synthesized and reacted with differ-
ent salts of Cu (II) ion (OAc−, NO3−, SO4 and Cl−) in absence and presence
2−
Correspondence
of secondary ligands (L0); 8-hydroxyquinoline, 1,10-phenanthroline or SCN−;
to form binary and ternary Cu(II)-chelates. The ligand and its Cu(II)-
complexes were fully characterized by analytical, spectral, thermal, conductiv-
ity and magnetic susceptibility measurements. The metal chelates showed
octahedral, square planar and /or distorted tetraherdal arrangements. Coats–
Redfern equations used to calculate the kinetic parameters of the thermal
decomposition stages (Ea, A, ΔH, ΔS and ΔG). The compounds exhibit lumi-
nescence property; promising interesting potential applications as photoactive
materials. Lippert–Mataga, Bakhshiev, Kawski–Chamma–Viallet and micro-
Omima M.I.Adly, Department of
Chemistry, Faculty of Education, Ain
Shams University, Roxy, Cairo, Egypt.
Email:omima_adly@yahoo.com
N
scopic solvent polarity parameter and ET correlation methods were applied
on the solvatochromic shifts of emission spectra to evaluate the ground (μg)
and excited (μe) states dipole moments. Excited state dipole moment is larger
than the ground state which may be attributed to π-π* transition. The coordi-
nating anions play an important role on the position and intensity of emission
band. The ligand and its metal complexes showed antimicrobial activity
towards Gram–positive bacteria, Gram–negative bacteria, yeast and fungus.
The molecular structural parameters of HL and its Cu(II)- complexes have
been calculated on the basis of DFT engaged in the Gaussian 09 program at
the B3LYP/6-31G(d,p) level; the theoretical data are correlated with the
experimental data.
K E Y W O R D S
DFT/B3LYB method, molecular-microscopic solvent polarity parameter, ONS hydrazone ligand;
fluorescence, Solvatochromic shift method
1 | INTRODUCTION
photochemistry and photophysical studies.[1,2] The effect
of solvents causes a solvatochromic shift that used to esti-
The effect of solvents on the absorption and fluorescence
properties of organic molecules and their complexes have
been a subject of several investigations in the field of
mate the electric dipole moment of short-lived species
which is a direct measure of the electron distribution in a
molecule. Dipole moments of the short-lived species are
Appl Organomet Chem. 2020;e5763.
wileyonlinelibrary.com/journal/aoc
© 2020 John Wiley & Sons, Ltd.
1 of 22
https://doi.org/10.1002/aoc.5763

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ADLY ET AL.
of considerable interest to give the important information
on the electronic, reactive centers and geometrical struc-
ture of the molecule.[3] Experimental data of the excited
state dipole moment is also useful in the semi-empirical
quantum calculation for short-lived states.[4] The previ-
ous knowledge of the dipole moments of the electroni-
cally excited species is often useful in the design of
nonlinear optical materials[5] and in the photochemical
transformation as well as in the elucidation of the nature
of the excited states.
the emission spectra of copper (II) complexes are investi-
gated to estimate the ground and excited states dipole
moments using solvatochromic shift methods. The den-
sity function theory (DFT) at 6-31G (d, p) level is used to
calculate the structural parameters of the optimized
structures. Antimicrobial activity of the hydrazone and
its copper complexes are monitored by agar disc diffusion
method against some species of pathogenic bacteria
(Staphylococcus aureus, Bacillus subtilis, Salmonella typ-
himurium and Escherichia coli), in addition to yeast (Can-
dida albicans) and fungus (Aspergillus fumigatus).
Various methods are available for the estimation of
the dipole moments and classified as either external or
internal. External methods include electro-optical
methods such as stark splitting or electric dichroism
which are known to give very accurate values of dipole
moments but need specialized equipment.[6,7] However,
the internal methods are solvatochromic shift manner
based on the spectral shifts caused as a result of the sol-
vent medium. This is a simple method used for the vari-
ety of molecules.[8,9] The changes in polarity (dielectric
constant) and polarizability (refractive index) of the sol-
vent can be attributed by the spectral shift produced by
the solvent. Copper (II) complexes play an important role
in the active sites of a large number of metalloproteins in
biological systems and potential application for numer-
ous catalytic processes in living organisms that involve
electron transfer reactions or activation of some anti-
tumor substances.[10] Copper (II) chelates exhibit
antibacterial, antifungal and anticancer activity.[11–13]
Chromones are common as a vital heterocyclic com-
pounds in organic synthesis, constitute one of the major
classes of naturally occurring and they are useful as bio-
logically active agents.[14–16]
2 | EXPERMENTAL
2.1 | Synthesis of 2-[(2-aminochromon-
3-yl)methylidene]-N-phenylhydrazine
carbothioamide (HL)
An ethanolic solution of N4-phenylthiosemicarbazide
(1.67 g, 10 mmol) was added dropwisely to an equimolar
amount of 2-aminochromone-3-carboxaldehyde (1.89 g,
10 mmol)[17] with continuous stirring. The reaction mix-
ture was heated under reflux for 15 min, a yellow crystals
seprated during reflux. The product was allowed to cool,
then filtered and rꢀecrystalized from ethanol; yield (2.47 g,
73%), m.p. > 300 C. The detailed synthetic steps for the
hydrazone ligand (HL) are given in details in Scheme 1.
2.2 | Synthesis of the copper(II)-
complexes
The present work describes the syntheses of a new
series of binary and ternary complexes of monoprotic tri-
dentate ONS-donor ligand, 2-[(2-aminochromon-3-yl)
methylidene]-N-phenylhydrazine carbothioamide (HL),
(Scheme 1), with various copper (II) salts, to investigate
the effect of anions, and presence of auxiliary ligand on
the structure and their spectral data. Effect of solvents on
Binary complexes were prepared by mixing a methanolic
solutions of HL ligand with a respective copper (II)salts
in molar ratio 1:1 then refluxed for 3 hr. However, the
ternary complexes were prepared by adding the second-
ary ligand (L0: 8-HQ, 1,10 phen or SCN−) to the nitrato
complex 2in a ratio 1:1:1 (M: HL: L0). The complexes
were filtered, washed with methanol then ether and dried
in a desicator over anhydrous CaCl2.
2.2.1 | Synthesis of binary [Cu(L)(OAc)
(H2O)] complex (1)
Cu(OAc)2.H2O (0.60 g, 3.0 mmol) dissolved in methanol
(30 ml) was added gradually to the solution of (HL)
ligand (1.01 g, 3.0 mmol). The reaction mixturewas
refluxed for 3 hr; dark green precipitate was obtained and
filtered off and washed several timeswith methanol,
diethylether and finally dried ina desicator over anhy-
drous CaCl2; yield (0.68 g, 51%).
SCHEME 1 HL ligand

ADLY ET AL.
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2.2.2 | Synthesis of [Cu(L)(SCN)], complex
(6)
without any further purification. In order to avoid the
effect of self-absorption and aggregation formation, the
measurements were taken in dilute solution (1 × 10−4 M).
Cu(NO3)2.2.5H2O (0.70 g, 3.0 mmol) dissolved in metha-
nol (30 ml) was added gradually to the solution of HL
ligand (1.01 g, 3.0 mmol) and heated for 30 min, then
KSCN (0.29 g, 3.0 mmol) in methanol (20 ml) was
addded. The reaction mixture was heated for 3 hr; green
precipitate was formed and filtered off and washed sev-
eral times with methanol, diethylether and finally dried
in a desicator over anhydrous CaCl2.; yield (0.86 g, 58%).
2.6 | Analyses and instrumentation
Elemental analyses were performed on a Perkin–
1
ElmerCHN-2400 analyzer. H NMR (400 MHz) and 13C
NMR (100 MHz) spectra were measured on Mercury-
300BB, using DMSO-d6 as a solvent and TMS (δ) as the
internal standard (Ministry of Defense, Chemical War
Department, Cairo, Egypt). The metal ion contents were
determined complexometrically according to literature
procedure.[20] Infrared spectra were recorded on FTIR
Nicolet IS10 spectrophotometer (cm−1), using KBr disks
at Central Laboratory, Ain shams University, Cairo,
Egypt. Mass spectra were obtained using GC-2010
Shimadzu Gas chromatography instrument mass spec-
trometer (70 eV). The Shimadzu TGA-50H instrument
was used to carry out thermal gravimetric analysis (TGꢀA)
of metal complexes from room temperature up to 800 C
with a heating rate of 10 ꢀC min−1(Micro-Analytical Cen-
ter, Cairo University, Egypt). The electronic spectra of
the ligand and its metal complexes were measured with a
2.3 | Theoretical methodology
All molecular calculations were performed in the gaseous
phase using density functional theory (DFT) with B3LYP
through exchange correlation functional.[6,7] The basis
set 6-31G (d, p) was used for all atoms using the GAUSS-
IAN 09 program package by means of Gauss View visuali-
zation program.[6,7]
2.4 | Antimicrobial screening
The antimicrobial activity of the tested compounds was
JASCO
model
V-550UV–VIS
spectrophotometer
determined using
a
modified agar disc diffusion
(200–900 nm). But, the fluorescence measurments were
recorded on perkin-Elmer650-10Sspectro-fluorimeter,
equipment with 1-cm quartz cell and 150 W Xenon Arc
lamp at room temperature. Magnetic susceptibility data
were measured using a Johnson Matthey, Alfa product,
Model No.(MKI) (Central Laboratory, Ain shams Univer-
sity, Cairo, Egypt) Finally, melting points were deter-
mined on a digital Stuart SMP3 apparatus.
method.[18,19] The activities were done at 500and
1000 μg/ml concentrations of ligand and its copper (II)
complexes in DMF solvent by using two pathogenic
Gram positive and Gram negative bacteria (Staphylococ-
cus aureus, Bacillus subtilis, Salmonella typhimurium and
Escherichia coli) and yeasts (Candida albicans) and fun-
gus (Aspergillus fumigatus).[18,19]
2.5 | Materials and methods
3 | RESULTS AND DISCUSSION
Copper(II)salts including acetate, nitrate, sulphate and
chloride were used. 2-Hydroxyacetophenone, phosphorus
The current HL ligand was allowed to react with sev-
eral Cu (II) salts; AcO−, NO3 , SO4 and Cl−; to inves-
tigate the role of counter anions on the structure and
spectral properties of products. Moreover, the ligand
was allowed to react with copper (II)nitrate in the pres-
ence of secondary ligands (L0) such as N-, N,O and N,N-
donor types; namely: SCN−, 1,10-phenanthrolin and
8-hydroxyquinoline, respectively. The synthesized com-
plexes are stable at room temperature, non-hygroscopic
and sparingly soluble in water and common organic sol-
vents. The present complexes are characterized by ele-
mental and thermal analyses, IR, UV–Vis and mass
spectra as well as conductivity and magnetic measure-
ments. Physical and analytical data of the prepared
complexes are listed in Table 1.
−
2−
oxychloride,
hydroxylamine
hydrochloride,
N4-
phenylthiosemicarbazide, were either BDH or Merck
products. 8-Hydroxyquinoline, potassium thiocyanate,
1,10-phenanthroline were either BDH or Merck products.
The selection of organic solvents for the present inves-
tigation was based on gradual change of their bulk polar-
ity (ε) and refractive index (n) parameters. The list of
solvents were methanol (MeOH), ethanol (EtOH),
propan-2-ol
(2-PrOH),
1,4-dioxane(1,4-Diox),
dimethylformamide (DMF), ethyl acetate (Etac), tetrahy-
drofuran (THF), benzene (benz), chloroform (CHCl3),
dichloromethane (DCM) and acetone (Me2CO). All of
these solvents were of spectroscopic grade and used

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ADLY ET AL.

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5 of 22
3.1 | FT-IR spectra
Table 2 depicts the characteristic FT-IR data and their
tentative assignments of the free ligand and its Cu (II)-
complexes. The formation of the current hydrazone
ligand is confirmed by the appearance of a new intense
vibrational band at 1605 cm−1 which assigned to the
azomethine (C=N) group, in addition the streching vibra-
tional bands appeared at 3343, 3219 (NH2, NH), 1640
(C=Oγ–pyrone) and 1264 (C=S) cm−1.[21]
To infer the coordination modes FT-IR spectral bands
of the complexes are compared with the analogous of the
free ligand, HL. All complexes showed a new broad band
in the range 3331–3469 cm−1 assignaed to ν (OH) of the
coordinated/uncoordinated water or methanolic mole-
cules with complexes.[21] The shift of the bands of the
free ligand at 1640 and 1605 cm−1 to a lower frequncies,
in all complexes, confirm the participation of (C=O) and
(C=N) functional groups in chelation.[21] This result was
further emphasized by the following relationships:
ν
C=C/cm−1 = −1313.3 + 1.7342 νC=O/cm−1, r = 0.95,
n = 5, except: HL, 3, 7 and 8; νC=C/cm−1 = −9920.1 + 7.149
C=N/cm−1, r = 0.96, n = 5, except complexes 2, 6, and 7.
ν
The positive slopes of these relationships, refer to the
dynamic electron re-distribution of the π – electrons over
the whole molecule to over-compensate the donation of
carbonyl and azomethine groups to the central Cu (II).
Moreover, the recompense of C=N is much higher (≈ 3.5
times) than C=O, as indicated from their slopes coeffi-
cients, in addition to a little elongation of C=N(18) com-
pared to C=O as shown in Table 12 see below (vide
infra).
The fading of thione band, ν(C=S) at 1264 cm−1 of
HL ligand in all complexes except complex 3, and
appears a new band at 690 cm−1 that assigned to ν(C-S),
providing an evidence to participtate as thiolate in chela-
tion (Scheme 2).[21] Further evidences could be extracted
from the tautomeric structures of copper (II) complex
dipected in Scheme 3, recommend a reduce of electron
densities at chelate centers (O,N). Different sources are
compensate this decrease in electron deficieny such as:
mesomeric effect of amino group at position 2 which
compensate the electron defeciencies at O and N centers
as illustrated in tautomers AandB, repectively. Beside,
repelling of the nitrogen atom that adjacent to the phenyl
ring and the sulfur atom (tautomers C and D). The latter
effect converts the sp2 - hybridization of non- chelated
nitrogen to sp3 hybridization. Concequently, the electron
density at this molecule zone is raised and recompense
the electron deficiency at the chelated centers, especially
at N center. This explains, the observed results of the
compensation at the chelated azomethine nitrogen about
three times as carbonyl oxygenatom.

6 of 22
ADLY ET AL.
metal bonds. Concequently, the postive slope of νCu-N
versus C=O, νCu-N/cm−1 = 1244.8 + 1.0408 νC=O/cm−1
,
r = 0.91, n = 5, except: complexes 1, 6and 8, reveals that
the strong interaction of C=N with copper (II)was
accompanied by high extent of back donation (spillover
of copper ion) to the carbonyl group (pile up). Extra evi-
dence, could be deduced from the positive slope of νC=N
SCHEME 2 Tautometic structures of HL
versus νC=O; νC=N/cm−1 = −764.93 + 1.4471 νC=O/cm−1
,
r = 0.98, n = 7, except: HL and complex 2, which points
again to the corresponding back donation at the
azomethine and carbonyl groups.
The coordinating modes of anions are inferred from
their IR spectra, the vibrational bands observed at 1435
and 1331 cm−1 of complex 1could be outstanding to
νas(COO−) and νs(COO−); with a separation, Δν = (νas-
νs) = 104 cm−1, designates bidentate character of the
acetato group.[21] Whereas, complexes 2, 7and 8displayed
new absorption bands at (1384, 1383 and 1384) and
(818,825 and 850) cm−1, respectively, indicating the ionic
3.2 | Mass spectra
The mass fragmentation patterns of hydrazone, HL,
ligand is depicted in Scheme 4 (Supplementary Figure 1),
showed the molecular ion peak at m/z 338that coincides
with the formula weight (338.39) proposed on the basis
of elemental analysis of HL ligand (Table 1). Moreover,
the mass spectra of complexes 1,4 and 5 (Supplementary
Figures 2–4) are selected as representative complexes,
showed the molecular ion peaks at m/z 477,436 and
738,respectively. This finding agrees with the formula
weights of their anhydrous complexes; [Cu(L)(OAc)
(H2O)].H2O (F.Wt = 495.99), [Cu(L)Cl] (F.Wt = 436.42)
and [Cu(L)2].2H2O (F.Wt = 774.37).
nature of the NO3 group.[21] While, the vibrational
−
bands observed at 1110 and 1067 cm−1 for complex
3could be assigned to monodentate sulfato group.[21]
Moreover, the shift of thione group band from 1264 cm−1
in free ligand, HL to 1242 cm−1 for complex 3,supports
chelating as thione instead of thiolate for other
complexes.
However, the new bands viewed at 1496 and
1495 cm−1 for mixed-ligand complexes 7and 8, respec-
tively, underneath the coordination of the C=N group of
the secondary ligands to the metal ion.[22] But, complex
6showed a characteristic bands at 2100 cm−1, medium
band at 746 cm−1 and weak band at 443 cm−1
corresponding to v(C ꢁ N), v(CS) and δ (NCS); respec-
tively, points to a unidentate character through nitrogen
atom of thiocyanato.[21] Finally, appearance of new weak
to medium intensities bands in the ranges: 527–585 and
403–470 cm−1 matched with ν(M-O) and ν(M-N);
respectively.[21]
1
3.3 | H-NMR spectrum of HL ligand
1
Figure 1 shows the H-NMR spectrum of HL ligand in
deuterated dimethylsulphoxide (DMSO-d6) in absence
and presence of D2O; the results are listed in Table 3. The
signal observed at δ 8.75 ppm assigned to the NH2 pro-
tons at position 2of chromone moiety. This signal dis-
appeared in the presence of D2O. The azomethine proton
was observed at δ 8.73 ppm. The two thioamidic NH pro-
tons in HL ligand display signals at δ 9.89 and
The assignments mentioned above are further empha-
sized by the linear regression analysis (LRA) between the
streching frequncies of coordinating centers and their
SCHEME 3 The possible tautomeric structures of
Cu(II)complex

ADLY ET AL.
7 of 22
dimethylsulphoxide (DMSO-d6) and data is listed in
Table 3. The spectrum showed 17signals, which is
equal to the number of carbon atoms in the HL
ligand. Characteristic signals appeared in the spectrum
at δ 91.8 ppm, this lower chemical shift value is due
to the electron repelling affinity of the amino group
that deshield C-3 position. However, the signals
observed at downfield region at δ 174.9 and 173.4 ppm
were attributed to C=Sthioamidic and C=Oγ–pyrone
,
respectively. Moreover, the signals attributed to C-2
and C-8a were observed at δ 161.7 and 152.5 ppm.
The higher chemical shift values might be attributed
to the electron withdrawing of mesomeric and induc-
tive effects of the carbonyl and nitrogen of the amino
group, and adjacent oxygen of the chromone moiety.
Finally, the six signals of the phenyl carbons appeared
in the range δ 125.2–133.3 ppm.
3.5 | Conductivity measurements
Molar conductivitiy data of the present complexes in
DMF solutions (10−3 M) measured at 25 1 C are col-
ꢀ
SCHEME 4 Mass fragmentation patterns of the HL ligand
lected in Table 4. The molar conductivity values of Cu
(II)-complexes can be classified into two categories: (a)
1:1 electrolytes for complexes 2, 7and 8have conductivity
values: 93,84 and 131 Ω−1 cm2 mol−1, respectively, and
(b) non-electrolytes for other complexes with conductiv-
ity values in the range 7–27 Ω−1 cm2 mol−1, except chloro
complex 4 that has 39 Ω−1 cm2 mol−1. The deviation of
complex 4 might arised from the partial replacment of
chloride ion by DMF solvent.[23] Thus, the conductivity
measurements of the copper (II)-chelates supports the
structure proposed in accordance with the results of ele-
mental analyses and IR spectra.
11.45 ppm; these signals vanished in the presence of
D2O.The doublet signal observed at δ 8.00 ppm is charac-
teristic for H-5 of chromone moity.
3.4 | 13C-NMR spectrum of HL ligand
Figure 2 demostrates the 13C-NMR spectra of the pre-
sent
hydrazone,
HL,
ligand
in
deuterated
FIGURE 1 1H-NMR spectrum of
HL ligand in (DMSO-d6) (A) in absent
of D2O and (B) in presence with D2O

8 of 22
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TABLE 3
1H and 13C NMR spectral data of HL ligand
Chemical shifts (δ, 1H)
Chemical shifts (δ, 13C)
in ppm (DMSO-d6, 300 MHz)
Assignment
in ppm (DMSO-d6, 75 MHz)
Assignment
C-3
7.17
(t, 1H, H-6, J = 7.8 Hz)
(m, 6H, Ph-H and H-8)
(t, 1H, H-7, J = 7.5 Hz)
(d, 1H, H-5, J = 7.2 Hz)
(s, 1H, CH=N)
91.8
116.4
121.4
123.5
124.8
125.2
126.6,
127.9,
128.2,
128.6
133.0
139.5
142.7
152. 5
161.7
173.4
174.9
7.33–7.47
7.69
C-8
C-4a
8.00
C-7
8.73
C-6
8.75
(br, 2H, NH2(a)
Ph-C
Ph-C
Ph-C
Ph-C
Ph-C
Ph-C
C-5
exchangeable with D2O)
(br, 1H, NH(b),
9.89
exchangeable with D2O)
(br, 1H, NH(c),
11.45
exchangeable with D2O)
CH=N
C-8a
C-2
C=O as C-4
C=S
3.6 | Electronic spectra and magnetic
measurements
the stability of the complex formation” versus the wave
number of LMCT, Eg/eV = 0.1332 + 9.665x10−6
ν
LMCT/cm−1, r = 0.94, n = 5, except 1, 5 and 7 complexes.
The electronic spectral data of HL ligand and its copper
(II) complexes recorded in DMF solution, Nujol mulls
and/or reflectance are listed in Table 4. The ligand has
two absorption bands at 289 nm (34602 cm−1) and
355 nm (28169 cm−1) assigned to the π-π* and n-π* transi-
tions of chromone ring and amide group, respectively.[24]
The intraligand transitions observed for the present
Electronic spectra of the current copper (II) com-
plexes might be either in D4h or C4v symmetry, for eg and
t2g levels of 2D free ion terms will split into B1g, A1g, B2g
and Eg levels, respectively. Thus, the three spin allowed
transitions expected for the visible and near IR regions.
All copper (II) complexes have broad d-d combination
bands in the range 680–532 nm (14706–18797 cm−1).[25]
These bands may be assigned to the following transitions:
chelates in the range 289–393 nm (34602–25445 cm−1
)
2
2
2
showing a red shift which might attributed to the dona-
tion of electron pairs of coordinating centers toCu(II)ion
on complexation. However, the high intense bands
observed in the region 407–453 nm (24570–22075 cm−1)
reveal ligand-to-metal charge transfer (LMCT) transitions
which might originated from O ! Cu and N ! Cu, tran-
sitions. The extent of these interactions can be deduced
from the postive slope of the linear relationship of
Egap(Egap = ELUMO - EHOMO, vide infra) “that measures
2B1g ! A1g (dx2 -y2 ! dz2), B1g ! B2g (dx2-y2 ! dxy)
and 2B1g ! Eg (dx2- y2 ! dxz, dyz) in the order of increas-
2
ing energy, according to the distortion degree resulted
from the ligand field and Jahn–Teller effect.[26] The elec-
tronic spectra of the present copper (II) complexes
showed one band around 600 nm corresponding to
2
2B1g ! B2g (dx2- y2 ! dxy) transition and this is consis-
tent with an octahedral configuration.[26] However, com-
plexes 3, 4and 6 displayed bands around 579,626 and

ADLY ET AL.
9 of 22
FIGURE 2 13C-NMR spectrum of HL ligand in (DMSO-d6)
TABLE 4
No.
Electronic spectra, magnetic moments and molar conductivity data of the HL ligand and its complexes
Electronic spectral bandsa (nm)
Conductancea
Complex
λ
max(DMF/nm) (Nujal mull) [refl]
μ
eff.B.M.
----
1.95
(Ω−1 cm2 mol−1
)
HL
289,355
----
7
1
2
3
4
5
6
7
8
[Cu(L)(OAc)(H2O)]. H2O
[Cu(L)(CH3OH)3]NO3
[Cu(L)(SO4)]
289,353,438, 680 [389,579,700sh]
291,353,432,640 [407,584]
293,393,422,579 [282,382,433,584]
350,437,626 [390,630]
1.95
1.75
1.79
1.96
1.76
2.31
1.88
93
17
39
13
27
84
131
[Cu(L)Cl]
[Cu(L2)].2H2O
290,353,407,532(472,592) [387,600]
274,351,413,614 [394,534,686]
289,353,437,648
[Cu(L) (SCN)]
[Cu(L)(8-HQ)(H2O)]NO3
[Cu(L)(Phen)(H2O)] NO3
353,453,670 (277,381,527) [393,434,670]
532 nm, proposing square planar and/or distorted
tetrahedralgometries. The magnetic moment values of
the current copper (II) complexes are found in the range
1.75–2.31 B.M., which are close to spin only value 1.73
B.M.corresponding to single unpaired electron of copper
(II)-d9 complexes.[27]
enhancement of fluorescence intenstity of all complexes
compared to the free ligand, except sulfato complex 3,
which may be attributed to the bridge connection of the
ligand to the metal ion, that superior its conformational
rigidity that diminishes the non-radiative energy
loss.[28,29] The emission band of HL ligand in DMF solu-
tion observed at 437.43 nm exhibits a red shift to the
range 498.46–585.77 nm for complexes upon excitation at
395and 436–477 nm, respectively. The effect of anion on
the emission spectra submits the following order: Cl−
3.7 | Fluorescence spectra
(492 nm) < SCN− (529 nm) < OAC− (539 nm) ≈ SO4
—
Table 5 collects the emission spectral data of the free
ligand, HL, and its binary and ternaray copper (II)com-
plexes (1–8) at room temperature in solid state and their
1 × 10−4 M DMF solutions. The data in Table 5 reveals
−
(540 nm) < NO3 (583 nm). This trend is in the opposite
dirction of anion's Lewis basicity (DNX) as shown in
Figure 3.[30] This empasized by the negative slope of the

10 of 22
ADLY ET AL.
TABLE 5
Florescence spectral data of HL ligand and Cu(II)-complexes in solid and DMF solution [1x10−4 M] at room temperature
Complexes
HL
λex/nm
395
395
464
444
465
445
458
428
332
362
436
386
457
477
467
437
477
437
λem/nm
437.43
486.23
538.73
583.50
540.48
585.77
540.48
569.37
492.36
487.18
498.46
434.81
529.41
572.52
546.00
575.11
532.18
575.94
Stoke Shift/nm
42.43
Intensity
Solution
Solid
4.4855 x106
5.4789x106
5.4753 x106
2.4928 x106
4.9971x106
5.0206x106
4.9112 x106
3.4872 x106
5.378 x106
91.23
(1)
(2)
(3)
(4)
(5)
(6)
(7)
(8)
[Cu(L)(OAc)(H2O)].H2O
[Cu(L)(CH3OH)3] NO3
[Cu (HL)(SO4)]
Solution
Solid
74.73
139.50
75.48
Solution
Solid
140.77
82.48
Solution
Solid
141.37
160.36
125.18
62.46
[Cu (L)Cl]
Solution
Solid
[Cu(L)2].2H2O
Solution
Solid
48.81
[Cu(L)(SCN)]
Solution
Solid
72.41
95.52
[Cu (L)(8-HQ)(H2O)]NO3
[Cu(L)(Phen)(H2O)]NO3
Solution
Solid
79.00
138.11
55.18
Solution
Solid
138.94
as shown in Table 5. This might recognize to the effect
of solute- solvent interaction on the emission spectra.
Moreover, the effect of complex stability on the range
of the Stoke shift of DMF solutions, 42.4–
160.4 nm = 118 nm, is more pronounced than in solid
state, 48.8–141.4 nm = 92.6 nm. This emphasized by
the negative slope, νStoke/(DMF)/cm−1 = 1.2185x105–
72.887νC=O/cm−1, r = 0. 95,n = 4, except: 2, 3, 4and 7,
this reveals high dependence of the Stoke shift on the
solute-solvent interaction and stability of complex.
However, Stoke shift in solid, νStoke//cm−1 (solid) = −
90004 + 58.459 νC=O/cm−1, r = 0.96, n = 4, except: 2,
4, 5and 7showed the opposite trend (positive slope),
indicating the main role of solute- solvent interaction
on the emission spectra. A similar manners observed
for bathochromic Stoke shift with the increasing of C-
S bond streching frequency, νC-S/cm−1 = −12264 + 22.109
FIGURE 3 Effect of anion on Cu(II)-complexes, λem/nm in
DMF vs DNx inDMF
linear relationship of λem versus the anion's DNX both in
DMF solution, λem/nm = 711.81–5.859 DNX, r = 0.99 and
n = 4. i.e.the higher basicity of anion's produced higher
stability complex that augments the conformational rig-
orousness of the electron mobility. So, blue shift of λem
with the increasing of anions DNX. This behavior con-
firmed by the negative slope of EHOMO “measures the
ionization potential of the compound” (vide infra) versus
the anion's DNx, EHOMO/eV = − 0.057 – 4.23x10−3 DNx,
r = 0.97, n = 4. Thus, the fluorescent emission may be
attributed to the intraligand fluorescence (π ! π*).[31]
The extent of Stoke Shift in solid state is higher
than their DMF solution, except for chloro complex 4
ν
, r = 0.94, n = 4.
Stoke/DMF
The spectral shift of absorption maxima wave number
(υa) and fluorescence maxima wave number (υf) along
with the Stokes shift (υa -υf) and an arithmetic mean of
wave numbers (υa + υf)/2 of the complex 5molecule in
different solvents are shown in Table 6. Complex 5was
chosen as an example to investigate the effect of solvent
on the absorption and emission spectra. The selection of
complex 5as a probe is based upon it's red shift of emis-
sion band from (λem. = 434.81 nm) in solid state to λem.
= 498.46 nm in DMF solution, which is in contrast to the

ADLY ET AL.
11 of 22
behavior of the rest complexes that exhibit a blue shift
(Table 4).
From Table 6, it is observed that emission maxima
wavelength are more pronounced than the absorption
maxima wavelength (little bite affected) for all the sol-
vents used. From the spectral shift data one can identify
the spectral transition namely n ! π* and/or π ! π*.
The data in Table 6 refer to the increase of Stoke's Shift
(red shift) accompanied with the increase of solvent
polarity suggesting π ! π* transition in the spectral levels
of the solute molecule.[32] So, the bands derivation from
π–π* transitions exhibit significant charge-transfer char-
acter resulting from electron donating part to the
electron-deficient part. The stabilized of the excited state
by solvent polarity is more pronouned than the ground
state of the molecule. As a result, the lowering of energy
for the HOMO to LUMO electron transition leads to a
bathochromic shift of the spectra with increasing solvent
polarity.
To characterize how a complex molecule interacts
with the surrounding species in solution, the emission
spectral data was correlated with the solvent parameters.
Furthermore, estimation of the dipole moments of a sol-
ute in the ground (μg) and excited (μe) states is also of
special interest to provide information about the change
in the electronic charge distribution upon excitation.
Consecutively, to estimate the ground state and excited
state dipole moments of complex 5, the solvent polarity
Lippert-Mattag, F1 (ε, n), Bakhshiev's, F2 (ε, n) and
Kawski-Chamma-Viallet's, F3 (ε, n) parameters and
Stokes shift (υa − υem) in addition to the arithmetic mean
of wave number (υa + υem)/2 were calculated as reported
in our previous work[22] (Table 6). Figures 4-6 (protic sol-
vents) show the plots of respective spectral shifts, Stokes
shift (υa − υem) and (υa + υem)/2 for complex 5 which
are observed in protic and aprotic solvents against the
FIGURE 4 The relation between νStoke/cm−1 vs F1

12 of 22
ADLY ET AL.
TABLE 7
Intercept (C), the slope (m), correlation coefficient
(r) and number of data points (n) corresponding to statistical
treatment of spectral shifts of Cu (II) complex 5, [Cu(L2)].2H2O
Eqn
Intercept (C)
1150.9
2180.9
1390.4
2402
Slope (m)
+5230.6 (m1)
4136.4
r
n
5
5
5
5
4
5
5
4
6
5
ν stoke vs F1
0.93
0.96
0.98
0.92
0.999
0.96
0.99
0.997
0.978
0.999
ν stoke vs F2
(νa + νf) vs F3
+1558.0 (m2)
1139.7
22247
−1329.4 (m3)
−2979.6
23492
n
FIGURE 5 The relation between νStoke/cm−1 vs F2
ν stoke vs ET
2134.2
1891.4
1877.7
337.21
+806.34 (m4)
+3409.4
ν stoke vs π*
+1482.2
+3283.3
as decribed previously[22] and collected in Table 8. The
Onsager cavity radius (a)of complex 5, found to be 8,88 Å
(Table 8) was calculated as described elswhere.[22] The
difference in dipole moment calculated from solvent per-
N
turbation method and the one calculated using ET are
reasonably in good agreement in case of protic solvents,
obviously indicating the excited state dipole moment to
be higher compared to ground state. The observed gro-
und state dipole moment of complex 5 is 6.71D, and
4.758 in protic solvents and 6.16 D and 0.743 in aprotic
solvents, respectively. However, the excited state dipole
moments are 12.41D and 15.194 D for protic group and
10.84 D and 9.667D.
The discrepancies observed between the experimental
results using diffrenet methods as well as the theortical
result possibly occur, in part from the limitation of the
Lippert-Mattage, Bakshiev and Kawski- Chamma-Viallet
equations used for the treatment of solvatochromic
effects. Indeed, a number of assumptions and simplifica-
tions are made concerning the validity of their use, such
as, (i) the dipole moments in the ground and excited
states are supposed to be approximately collinear[33] and
(ii) the solute molecules are considered to be spherical
and consequently, the dipole moment is supposed to
remain constant in the Onsager cavity.[34]
The comparatively higher value of μe than μg shows
the existence of difference in the electronic charge distri-
bution in the excited singlet state relative to the ground
state. This result confirms that an extended π electronic
delocalized system is present in the investigated complex,
with important charged resonance structures in the
excited singlet state. The π ! π* may increase the magni-
tude of charges leading to increase in the dipole moment
of complex 5 in excited state.
FIGURE 6 The relation between νStoke/cm−1 vs ET
N
polarity functions (F1, F2 and ETN) and F3, respectively. A
linear regression was done and the data were fit to
straight lines for all plots using the Kgraph 3 software.
The slopes (m), intercepts (C), correlation coefficients
and the number of data points of fitted lines for complex
5are summerized in Table 7.
The data in Table 7 demonstrates two groups: group1,
comprise protic solvents: CHCl3, 2-PrOH, EtOH, MeOH;
and group 2,includes aprotic solvents: Diox, Etac, DCM,
Me2CO, DMF. The interaction with weak polar solvents
depends on the dipole-induced-dipole forces, while with
aprotic (polar but non- hydrogen bond donar (HBD)) sol-
vents, the solute-solvent interaction depends on the
stronger dipole–dipole forces. In protic solvents (polar
but hydrogen bond donar (HBD)), in addition to dipole–
dipole interaction, specific interaction such as H-bonding
may be effective as the intermolecular charge transfer
(ICT) character is favorable for H-bonding with hydroxyl
groups present in protic solvents. Hydrogen bonding
interaction usually puts a severe restriction on the valid-
ity of the solvatochromic methods.
The slopes given in Table 7 are used to estimate the
ground (μg) and the excited (μe) states dipole moments

ADLY ET AL.
13 of 22
TABLE 8
The Onsager cavity radius (a), Ground-state and singlet excited state dipole moments (in Debye, D) for Cu(II) complex 5,
[Cu(L2)].2H2O
aμg
bμg
6.71
6.16
bμe
cμg
4.76
0.74
cμ
dΔμ
4.12
8.47
μe/μg
1.848
1.760
e
Complex 5
gp1
a/Å
8.88
9.28
12.41
10.84
15.19
9.67
gp2
gp1 2,3,7,9,10 (protic)
gp2 1,4,6,8.11 (aprotic)
aGround state dipole moments estimated by Gaussian09 using DFT software.
bcalculated using (F1, F2) equations.
ccalculated from Kawski-Chamma-Viallet's (F2, F3) equations.
N
dcalculated from ET
.
eusing F1, F2 results.
The change in dipole moment (Δμ) was calculated
also from the solvatochromic shift and molecular-
microscopic solvent polarity parameter methods which
are found to be 4.12 and 8.47 D for protic and aprotic
groups, respectively, as shown in Table 8. It is evident
from Table 8 that the ground state dipole moments
evaluated from Salvatochromic shift method (6.712 D)
is smaller than that from abinitio calculations using
DFT (9.278) for complex 5 probe. This difference could
be partly due to the simplification used in
solvatochromic methods and to strong specific effects of
the solvents of different nature. However, abinitio calcu-
lations gives dipole moment values only for molecule in
a gas phase. Numerous reports in literature have shown
disagreement between theoretical and experimental
values.[22,35]
increase of the Stokes Shift with solvent polarity reveals
an increase in the dipole moment on excitation. Probably
result,
A systematic solvent effect analysis is therefore, use-
ful to understand the effect of the environment and var-
ious mechanisms of de-excitation and relaxation of an
excited molecule in solution. Solvent dependence of the
emission spectra was observed for complex 5, it is
exhibited a red shifted emission from 483.4 to 507.6 nm
as the polarity of the solvent commonly increased with
some exceptions. Quantitative analysis to specific and
non-specific interactions using Kamlet–Taft method and
other solvent parameters such as Lewis acidity (acceptor
number, AN) and Rechardt's polarity (ETN) reveals two
main groups of solvents. The emission spectra of com-
plex 5solutions in THF, MeOH, Benz, 2-PrOH and
EtOH solvents displayed a red shift as indicated from
the negative slope for the linear solvation energy rela-
tionship (LSER) of νem versus the hydrogen bond donor
(HBD), α: νem/cm−1 = 20680–478.76α, r = 0.99, n = 5.
Whereas, in Etac, DCM, 2-PrOH, EtOH and DMF
If dipole moments μe and μg are form an angle φ,
since they are not parallel to each other. So, the angle ϕ
is calculated using the Equation given below that found
to be, 121.63 and 179.30, for protic and aprotic solvents
groups.
exhibits
em/cm−1 = 19752 + 707.9α, r = 0.99, n = 5, which is
1.5 times of the red shift.
Nevertheless, the hydrogen bond acceptor (HBA), β
showed a negative slope (red shift) in AC, Diox, Benz,
DMF solvents, νem/cm−1 = 20780–1565.1β, r = 0.92,
n = 4, with an extent of 2.6 times of the blue shift in
Diox, DCM, MeOH, 2-PrOH, EtOH solvents, as it has a
positive slope, νem/cm−1 = 19815 + 599.0 β, r = 0.97,
n = 5.
a
blue
shift
(positive
slope),
ꢂ
ꢃ
ꢀ
ꢁ
ꢀ
ꢁ
1
m3
m2
ν
2
2
2
cosϕ =
μg2 + μe
−
μg −μe
2μgμe
In addition to, the excited state is greater than the ground
state dipole moment. Thus, the variation in dipole
moment value can be explained in terms of possible reso-
nance structures of the investigated compound that
arised from the delocalization of π electrons.
The magnitudes of Stokes shift of complex 5are vary
between 2248.122–3233.675 cm−1 in THF and DMF solu-
tions. On changing solvents from a low polar to a high
polar it is notice a difference in Stokes Shift of about
985.56 cm−1, indicative again of charge transfer transi-
tion. The large magnitude of Stokes Shift showed that the
excited state geometry should be more stabilized by sol-
vent polarity than the ground state. Commonly, the
However, the non-specific effect of solvent polariz-
ability (π*) on the position of emission band,
ν
shift in THF, acetone, DCM, Benz, 2-PrOH, DMF and
blue shift (positive slope), νem/cm−1 = 19470 + 1111.2 π*,
r = 0.97, n = 5 in CHCl3, acetone, Etac, Benz, Diox,
MeOH solvents. Thus, the first group of solvent indicates
that the π* (non-specific) has doublet effect of β and α
em/cm−1 = 22599–3283.4 π*, r = 0.99, n = 5, reveals red

14 of 22
ADLY ET AL.
(specific) parameters of the solvents. Which are responsi-
ble for the photophysics properties of the present probe
in the ground state with a ratio 2: 1: 1.2, while the excited
state is mainly controlled by π* and β than α as infered
from their coeffient ratios 7: 3: 1(π*: β: α).
3.8 | Thermogravimetric analysis
Thermogravimetric analysis (TGA) are carried out for
complexes 1, 2, 6 and 8 as representative examples for
thermal analysis (Supplementary Figures 5–8), the results
of thermal analysis of these complexes are sumerized in
Table 9, which found to be in agreement with elemental
analyses. Thermodynamic and kinetic parameters have
been calculated using Coats-Redfern equations.[36]
The first decomposition stage of the complex 1extends
up to 133 ꢀC, related to the loss of associated water mole-
cules. The second stage of decomposition extends up to
283 ꢀC related to the loss of coordinated water, HCN and
HSCN molucules or methanol molecules. In case of com-
plexes 2and 7, the decomposition begins at 386and
366 ꢀC related to the loss of one HNO3 molecule, while in
complex 1the decomposition extends up to 317related to
the loss of one CH3COOH.
The kinetic and thermodynamic parameters are cal-
culated by using the steps of the thermal pattern of
decomposition of the ligand and its Cu (II) complexes. By
using the Coats-Redfern's method the activation energy
Ea, the pre-exponential factor A, enthalpy change (ΔH),
entropy of activation (ΔS), and the Gibbs free energy
change (ΔG) are estimated from TG curve and listed in
Table 10.[36,37]
Moreover, the specfic effect of solvent through solvent
polarity ETN, and Lewis Acidity (AN) parameters display
the following linear free energy relationships (LFER),
ν
em/cm−1 = 21562–24.86 ETN, r = 0.98, n = 6, red shift
for the excited stat in THF, Etac, MeOH, Benz, 2-PrOH,
EtOH solvents which is higher than the blue shift (posi-
tive slope), νem/cm−1 = 19322 + 19.93 ETN, r = 0.94,
n = 5, in CHCl3, Etac, 1,4-diox, 2-PrOH, EtOH solvents
for the ground state. Whereas, the effect of AN on the
ground state has a doublet effect than the excited state,
ν
em/cm−1 = 20795–14,17 AN, r = 0.99, n = 5, in THF,
MeOH, Benz, 2-PrOH, EtOH solvents and
em/cm−1 = 19352 + 27.98 AN, r = 0.93, n = 5, in CHCl3,
DCM, 2-PrOH, EtOH, DMF solvents.
ν
A general observation is the increase in Stokes' Shift
values with increase in solvent polarity indicating that
there is increase in the dipole moment upon excitation.
In such cases, the relaxed excited state S1 will be energet-
ically stabilized relative to the ground state S0 and hence,
a significant red shift of the fluorescence will be
observed.
TABLE 9
Compound
Thermogravimetric analysis of Cu(II) complexes of HL ligand
DTG peak
(ꢀC)
Temperature
range (ꢀC)
Decomposition product lost
(formula weight)
Weight loss (%) found
(calculated)
[Cu(L)(OAc)(H2O)].
H2O (1)
101
50–133
227–283
283–317
414–480
475–734
- H2O
4.20 (3.63)
8.40 (9.07)
266
- H2O+ HCN
CH3COOH + HSCN
-N2H2
308
23.40 (23.99)
6.15 (6.05)
439
554
C15H10O−
-CuO
42.13 (41.58)
15.72 (16.04)
5.90 (5.72)
Residue
178
[Cu(L)(CH3OH)3]
NO3(2)
110–208
209–282
282–386
387–538
-CH3OH
257
- 2CH3OH
-HNO3 + HCN
C16H11N3SO-
-CuO
11.90 (11.45)
15.79 (16.09)
51.91 (52.44)
14.65 (14.22)
25.13 (25.71)
6.30 (5.88)
366
467
Residue
247
[Cu(L)(SCN)] (6)
147–290
290–406
406–571
- 2HSCN
329
- HCN
498
-C15H10N2O
- CuO
51.46 (50.97)
17.11 (17.33)
14.72 (15.73)
15.52 (14.06)
58.52 (58.08)
11.24 (12.03)
Residue
228
[Cu(L)(H2O)(Phen)]
NO3(8)
87–260
260–366
366–680
- H2O + HCN + HSCN
- HNO3 + N2H4
-C27H16N2O
-CuO
298
527
Residue

ADLY ET AL.
15 of 22
TABLE 10
Compound
Temperatures of decomposition and the kinetic parameters of Cu (II) complexes of HL ligand
n
Δ E
ΔH
ΔS (kJ mol1
K−1
ΔG
Step order
T (K)
373.71
578.02
641.09
711.68
A (S−1
)
(kJ mol−1
)
(kJ mol−1
)
)
(kJ mol−1
94.78
)
[Cu(L)(OAc)(H2O)].
H2O (1)
1st
1
1.1726x107
53.12
21.12
82.88
18.27
51.99
50.01
16.31
77.55
12.35
45.12
131.49
48.06
17.88
136.79
27.54
105.75
91.13
27.05
129.19
−0.119
−0.261
−0.126
−0.257
−0.216
−0.044
−0.173
−0.256
−0.003
−0.219
−0.123
−0.074
−0.222
−0.102
2 nd
3 rd
4 th
5 th
1st
0
0.7274
167.43
158.85
195.86
224.32
111.37
139.66
206.97
134.82
161.08
200.68
127.39
159.82
210.52
0.33
0.5
1
8.676x 106
1.37
826.97 225.90
[Cu(L)(CH3OH)3]NO3
(2)
1
450.66
529.91
739.72
518.79
2.0351x1015 135.23
2 nd
4 th
1st
0.33
0
2.83 x 104
52.47
24.02
141.10
32.59
112.13
95.16
32.01
135.8
1.870
[Cu(L)(SCN)] (3)
1
4.66x1013
2 nd
3 rd
1st
1
607.27 112.79
1
767.52
484.90
596.80
800.93
1.512x107
[Cu(L)(Phen)(H2O)]
NO3 (4)
1
3.434x109
2 nd
3 rd
0.66
1
81.36
2.2639x108
From the data in Table 10, the kinetic and thermody-
namic parameters indicating that, The activation energies
of decomposition are found to be in the range 18.27–
141.10 kJ mol−1, the more positive values of the activa-
tion energy indicating that the thermal stabilty of the
compounds. The positive values of ΔH indicates that the
decomposation is endothermic, while the negative values
of activation entropies ΔS indicates that the complexes
are more activated than reactant. The positive values of
ΔH and ΔG indicating that non-spontaneous processes.
Finally, the structures of the present complexes
(Scheme 5) are proposed based on the tools mentioned
above.
occupied molecular orbital (LUMO) that could act as the
electron accepter. Energy difference between HOMO and
LUMO orbitals is called as energy gap (Egap) which deter-
mines the kinetic stability and chemical reactivity of the
molecule. Additional parameters, such as absolute elec-
tronegativities (χ), chemical potentials (I), absolute hard-
ness (ƞ), absolute softness (S), global electrophilicity (ω)
have been calculated.[22]
In Table 11, the energy gap (Egap/eV) of the present
compounds are found in the range 0.615–3.976 eV; the
most reactive is complex 7 in which the nergy gap is the
smallest value (0.615 eV). This can certainly facilitates
the charge transfer within the molecule than in the free
ligand HL (Egap = 3.976 eV).[39] In consequence, it is
assumed to be the softest complex 1though it displays the
low value of hardness (0.307); so is the most reactive
complex. However, the complex 1 shows the highest
energy gap with (2.751 eV). Therefore, it is admitted to be
the hardest complex (1.375 eV); it is the least reactive
complex. Among the most important trends of the ele-
ctronegativies (χ/eV) of the present compounds found in
the range (see Table 11) 2.699–4.595 eV, the lower elec-
4 | MOLECULAR ORBITAL
CALCULATIONS
All geometrical optimizations of HL ligand and its com-
plexes were performed with DFT method using
B3LYP/6-311G (pd) level implemented on the Gaussian
09.[38] The calculated structural parameters: dipole
moment, total energy, HOMO, and LUMO energies are
collected in Table 11 and shown in Figures 7-10.
Frontier molecular orbitals (FMO's) illustrated the
important concepts such as conjugation, aromaticity and
lone pairs. The hardness and softness of the molecule is
decide by FMO energies. Ionization potential of the mole-
cule is determined from the energy of the highest occu-
pied molecular orbital containing electrons (HOMO) that
could act as an electron donor. The energy of the lowest
,
tronegative is complex 2 [Cu(L)(CH3OH)3]NO3 which is
the cationic complex. However, the highest
elctronegative is complex 4, [Cu(L)Cl], which includes
the chloride ion that has the more electronegative ion.
The general trend of electronegativities (Table 11)
showed a good correlation with chemical intuition and
electronegativity anion.
However, the global electrophilicity (ω/eV) values
that estimated by using the electronegativity and

16 of 22
ADLY ET AL.
SCHEME 5 Representative structures of the Cu (II)
complexes 1–8
chemical hardness parameters (Table 11) are found in
the range: 2.808–29.404 eV.[22] Since, the global electro-
philicity (ω/eV) is a useful reactivity descriptor that can
be used to compare the electron-donating abilities of mol-
ecules.[40] i.e., A high value of electrophilicity describes a
good electrophile while a small value of electrophilicity
describes a good nucleophile.[41] So, the poor electrophile
of the current complexes is the complex 2, while the good
electrophile is complex 7.
inversaly directed with their Stoke's Shift, EHOMO/
eV = −0.116 – 9.21x10−6
n = 3.
ν
stoke/cm−1(solid), r = 0.99,
The polarity or charge separation over the molecule
of the present complexes, which is measured by the
dipole moment (μ) (Table 11), is found in the range
4.58–15.73 D, the less polar one is complex 2 but the-
highest polar is complex 8. Moreover, the positiveslope of
the linearship of the the ground state dipole moment ver-
sus the DNx, μ/D = − 4.745 + 0.476 DNx, r = 0.91, n = 4,
refers to the direct proportional of complex polarity with
anion's Lewis basicity.
The selected bond lengths extracted from the opti-
mized structures are reported in Table 12. The bond
lengths of the free ligand, HL, in the vicinity of coordi-
nating centers are generally more or less elongated for
most complexes. The theortical data in Table 12 reveals
the following elongation:
Application of the linear regression analysis to corre-
late the theortical with the experimental data. Since, the
relationship of the calculated Egap values versus the
experimental data supports the interpretaion given
above; in addition to realistic theortical data obtained.
The positive slopes of Egap that analogous to the stability
of complexes versus the stretching frequencies of vCu-O
and vCu-N those analogous to the extent of ligand –metal
interactions: Egap/eV = −0.68845 + 0.0013521 vCu-
O/cm−1, r = 0.99, n = 6, except: complex 5. Egap
/
ev = −0.78016 + 0.00187 vCu-N/cm−1, r = 0.99, n = 5,
except: 1 and 3; supporting the IR assignments. However,
the effect of the secondray ligands (HL, 8-HQ and
a Elongation of C-O15 (1.259 of HL) to become in the
range 1.268 to 1.318 Å.
b Elongation of C-N18 (1.309 of HL) to become in the
range 1.313 to 1.409 Å.
1,10-phen) of ternary complexes, showed that EHOMO
,

ADLY ET AL.
17 of 22
c Elongation of C-S24 (1.718 of HL) to become in the
range 1.754 to 1.838 Å.
This elonagation observed at coordinating centers
confirmed the coordination of HL with copper ion via C-
O15, C-N18 and C-S24. Tautomeric structures showed in
Scheme 2 point to the migration of the lone pair on the
nitrogen of NH2 group at position 2 to the chromone ring
(intramolecular charge transfer). This migration also
elongate C=O γ-pyrone, C=N
and the
chelated azomethine
C2 = C3 chromone, as listed in Table 12.
However, the shortness of CN19 (1.387 A of HL) to
become 1.313–1.362 Å in all complexes except complex 3.
The almost unchange of C-N19 of complex 3 suggest the
coordinated of C-S as thione instead of thiolat in other
complexes. This explain why the IR vibrational band of
C=S group in complex 3 is shifted from 1264 to 1242 cm−1
upon complexation with CuSO4. (see IR section). Conse-
quently, thione $ thiol tautmerism, through the migra-
tion of the proton on HN19 to the thione group. More
evidence, the negative slpoe of νC-Svs νC=O; νC-
S/cm−1 = 1120.6–0.2657 νC=O/cm−1, r = 0.96, n = 5, except:
HL, 2, 4 and 8; refers to electron redistribution (dynamic)
on the whole molecule during the complexation.
The formal charge of Cu-ion is +2, whereas, the com-
puted natural charges of the current complexes signifi-
cantly reduced to the range: +0.187 to +0.805 as shown
in Table 12. This reveals electron transfer of the mono-
anionic ligand, L ! Cu (II) as expected. This charge
transfer (CT) prototype is confirmed by the changes of
atomic charges on the ligand's donor atoms. The negative
charge value at N18 donor atom of free ligand,
HL − 0.338, experience a reduction to the range − 0.252
to −0.332, except complexes 2, 5 and 8 that exhibited
increment the negative charge on N18 (L Cu).
Whereas the charge at O15 donor atom of HL (−0.457)
increased in range − 0.463 to −0.700 for all complexes,
except complex 5that showed −0.33. Nevertheless, S24
donor atom in uncomplexed ligand (natural charge is
−0.062) reduces the negative charges ranges from −0.099
to +0.619 in all complexes.
Finally, it is wondered to see the reduction of the nat-
ural charge of non-coordinated center N19 of HL ligand
(Table 12) from −0.525 to the range − 0.162 to −0.514, in
all complexes, except complex 2which almost un-chan-
ged. This is an extra verification to resonance structures
that proposed in Scheme 2.
5 | ANTIMICROBIAL STUDIES
The results of antimicrobial assessment are listed in
Table 13. The ligand exhibits low activity against all

18 of 22
ADLY ET AL.
TABLE 12
The selected bond lengths and charge of Cu(II)complexes of HL ligand
Charges
Bond Length
Compound
O15
S24
N18
N19
Cu(II)
-
C=O15
1.259
1.280
1.299
1.284
1.300
1.295
1.280
1.268
1.318
C=S24
1.718
1.798
1.800
1.773
1.789
1.815
1.754
1.838
1.803
C=N18
C=N19
1.387
1.332
1.361
1.372
1.362
1.321
1.340
1.325
1.313
N-N
HL
1
−0.457
−0.544
−0.602
−0.609
−0.571
−0.330
−0.579
−0.463
−0.700
−0.062
0
−0.338
−0.332
−0.375
−0.300
−0.252
−0.369
−0.247
0.267
−0.525
−0.329
−0.531
−0.514
−0.332
−0.192
−0.300
−0.162
−0.297
1.309
1.313
1.409
1.307
1.345
1.342
1.346
1.287
1.298
1.412
1.389
1.349
1.419
1.320
1.367
1.406
1.348
1.395
0.560
0.621
0.805
0.519
0.187
0.294
0.698
0.262
2
−0.090
−0.099
0.060
0.619
0.255
0.429
0.464
3
4
5
6
7
8
−0.453
The increased activity for Cu(II)-chelates can be
explained on the basis of the chelation theory. It is recog-
nized that chelation improve the potent bactericidal
agent and widespread killing of microbes than the the
free ligand. It is convenient that, the positive charge of
the metal ion is partially shared with the donor atoms
present in the ligands and there by π-electron delocaliza-
tion over the whole chelation. This enhances the lipo-
philic character that favors its penetration the lipoid
layer of the bacterial membranes.[42]
DFT calculations may be correlated to the antimi-
crobial action of the studied compounds. The dipole
moments, which may give some insight on the degree
of hydrophobicity/hydrophilicity of the compounds
properties are of interest in this work. The data in
Tables 8 and 11 reveal inverse correlation between the
dipole moment and the activity of the isolated -Cu(II)-
complexes towards the studied bacterial and fungal
species. i.e, as dipole moment (polarity) decreases, such
FIGURE 7 Molecular modeling ofHL
organisms. The activity of metal complexes varies from
lower to intermedate activity, except complex 2that has
high activity towards Staphylococcus aureus (Table 13).
FIGURE 8 The electron density of HOMO and LUMO of HL ligand

ADLY ET AL.
19 of 22
FIGURE 9 Molecular modeling of Cu(II)complex1
FIGURE 10 The electron density of HOMO and LUMO of Cu(II)complex1
as complexes 2(4.58 D) and 7 (3.88 D). Hence, the
lipophilic nature of the compound increases, which
favors its penetration through the lipid layer of the
microorganism.[43] This in turn deactivates enzymes
responsible for respiration processes of the investigated
micro-organisms. i.e.the development of a lipophilic
nature could enhance its diffusion through the cyto-
plasmic membrane, and accordingly increase the
cellular uptake of metal ions by bacterial cells as
shown in Table 13.[43]
6 | CONCLUSIONS
The present work reports the effect of anions and second-
ary ligands on the geometrical structures of synthesied

20 of 22
ADLY ET AL.

ADLY ET AL.
21 of 22
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The effect of solvents on the absorption and fluores-
cence spectra of complex 5in different solvents are
reported at room temperature. The ground (μg) and the
excited state (μe) dipole moments of complex 5 were esti-
mated from Lippert's, Bakshiev's and Kawski-Chamma-
Viallet's equations using the solvatochromic shift methods.
The ground state dipole moment (μg) was also estimated
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the solvent and solute molecule on the π ! π* transitions.
The antimicrobial study reveals that the metal complexes
more active against Gram-positive than Gram-negative
bacteria. Studies suggested that there is an inverse correla-
tion between the dipole moment and the activity of the
complexes against the studied bacterial and fungal species.
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ORCID
Omima M.I. Adly https://orcid.org/0000-0003-0197-
8118
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Coordination Compounds, Sixth ed., Wiley, NewYork 2009.
[22] O.M.I. Adly, A. Taha, S.A. Fahmy, J.Mol. Struct. 2019,
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SUPPORTING INFORMATION
Additional supporting information may be found online
in the Supporting Information section at the end of this
article.
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How to cite this article: Adly OMI, Taha A,
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calculations and biological studies of
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