Synthesis, characterization, antimicrobial, and genotoxicity activities of acetoacetanilide-4-ethyl thiosemicarbazone complexes

Article abstract of DOI:10.1080/10426507.2010.503207

Chemical Equation Presented Co2+, Ni2+, Cu2+, Pd2+, Cd2+, UO22+, and VO2+ ions were found to form stable complexes with acetoacetanilide-4-ethyl thiosemicarbazone (H2L). The structures of the complexes were elucidated by analysis (elemental and thermal),

Full text of DOI:10.1080/10426507.2010.503207

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Synthesis, Characterization,
Antimicrobial, and Genotoxicity
Activities of Acetoacetanilide-4-ethyl
Thiosemicarbazone Complexes
R. R. Zaky a
a Department of Chemistry, Faculty of Science , Mansoura
University , Mansoura, Egypt
Published online: 19 Feb 2011.
To cite this article: R. R. Zaky (2011) Synthesis, Characterization, Antimicrobial, and Genotoxicity
Activities of Acetoacetanilide-4-ethyl Thiosemicarbazone Complexes, Phosphorus, Sulfur, and Silicon
and the Related Elements, 186:2, 365-380, DOI: 10.1080/10426507.2010.503207
To link to this article: http://dx.doi.org/10.1080/10426507.2010.503207
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Phosphorus, Sulfur, and Silicon, 186:365–380, 2011
Copyright ꢀC Taylor & Francis Group, LLC
ISSN: 1042-6507 print / 1563-5325 online
DOI: 10.1080/10426507.2010.503207
SYNTHESIS, CHARACTERIZATION, ANTIMICROBIAL,
AND GENOTOXICITY ACTIVITIES OF
ACETOACETANILIDE-4-ETHYL THIOSEMICARBAZONE
COMPLEXES
R. R. Zaky
Department of Chemistry, Faculty of Science, Mansoura University,
Mansoura, Egypt
GRAPHICAL ABSTRACT
O
NHC2H5
NHC2H5
S
PhHN
H3
CH3
HN
X
HN
N
S
C
N
Cl
CH3
N
O
S
HN
NHC2H5
Y
O
M
Cu
PhHN
S
Cl
O
O
O
H3
C
O
O
O
Cu
OH2
X
NHPh
O
O
CH3
OH2
[Cu(H2L)(OAc)2(H2O)].2H2O
[Cu(H2L)Cl2(H2O)2]H2O, [Pd(H2L)Cl2]
[Cu(H2L)(SO4)(H2O)].H2O
S
H
2
N
1
N
3
N
H
4
N
H
C2H5
O
CH3
H2
L
OH2
Cd
NHC2H5
CH3
H2
O
O
HN
S
V
N
O
O
S
2H2
O
O
M
S
NHPh
H2O
PhHN
N
CH3
C2H5HN
O
O
NHPh
N
N
C2H5HN
N
OH2
H3C
H3
C
[Ni(HL)(OAc)(H2O)].2H2O,
[Co(HL)(OAc)(H2O)].2H2O
[VO(L)(H2O)].2H2O
[Cd(L)(H2O)].H2O
Abstract Co2+, Ni2+, Cu2+, Pd2+, Cd2+, UO22+, and VO2+ ions were found to form stable
complexes with acetoacetanilide-4-ethyl thiosemicarbazone (H2L). The structures of the com-
plexes were elucidated by analysis (elemental and thermal), spectroscopy (electronic, IR, ESR,
and 1H NMR spectra), and physical measurements (magnetic susceptibility and molar conduc-
tance). IR spectra suggest that H2L acts as a monodentate, bidentate, and/or tridentate ligand.
The electronic spectra of the complexes and their magnetic moments provide information about
the geometries. The room temperature solid state ESR spectra of the Cu(II) complexes show
dx2-y2 as the ground state, suggesting tetragonally distorted octahedral geometry around the
Cu(II) center. The kinetic and thermodynamic parameters for the different decomposition steps
in [Ni(HL)(OAc)(H2O)] · 2 H2O and [VO(L)(H2O)] · 2 H2O complexes were calculated us-
ing the Coats–Redfern equation. Finally, the antimicrobial and genotoxic activities have been
tested.
Received 17 March 2010; accepted 20 May 2010.
Address correspondence to R. R. Zaky, Department of Chemistry, Faculty of Science, Mansoura University,
Mansoura, Egypt. E-mail: rania.zaky@yahoo.com
365

366
R. R. ZAKY
Supplemental materials are available for this article. Go to the publisher’s online edition of
Phosphorus, Sulfur, and Silicon and the Related Elements to view the free supplemental file.
Keywords Antimicrobial activity; genotoxic activity; ESR; thermal analysis; thiosemicar-
bazone complexes
INTRODUCTION
Thiosemicarbazones have been extensively studied because they have a wide range
of actual or potential medical applications,1 which include notably antiparasital,2 antibac-
terial,3 and antitumor activities.4,5 Many thiosemicarbazones, such as marboran or triapine,
are already used in medical practice. Their mechanism of action is still controversial in
many respects, but it is known that heterocyclic thiosemicarbazones act by inhibiting ri-
bonucleotide reductase, a key enzyme in the biosynthesis of DNA precursors.6–9 In general,
thiosemicarbazones are obtained by condensation of the corresponding thiosemicarbazide
with aldehydes or ketones. Thiosemicarbazones possess a wide range of biological activity,
depending on the parent aldehyde or ketone. Earlier reports on N(4)-substituted thiosemi-
carbazones have concluded that the presence of bulky groups at the N(4) position of the
thiosemicarbazone moiety greatly enhances biological activity.10–12
The present work aims to synthesize and characterize complexes of acetoacetanilide-
4-ethyl thiosemicarbazone (H2L), with Co2+, Ni2+, Cu2+, Pd2+, Cd2+, UO22+, and VO2+
ions. The possible modes of chelation, the geometry, and the nature of bonding of the
complexes are discussed on the basis of various spectroscopic methods (1H NMR, IR, UV-
Vis, ESR). In addition, the kinetic and thermodynamic characteristics of the decomposition
steps for Ni2+ and VO2+ complexes have been studied employing Coats–Redfern equations.
Their antimicrobial and genotoxic activities also have been tested.
RESULTS AND DISSUCTION
The isolated solid complexes are stable in air and easily soluble only in DMF and
DMSO. Most of the complexes decompose when heated at >300◦C. The molar conduc-
tance values in DMSO are 3–6 ꢀ−1 cm2mol−1, indicating that all complexes are non-
electrolytes.13 The color, melting points, and elemental analyses of the isolated complexes
are listed in Table 1.
IR and 1H NMR Spectra
The IR spectrum of H2L (Figure 1) shows three bands at 3328, 3249, and 3188 cm−1
assigned to υ(N4H), υ(N3H), and υ(N1H) vibrations (Figure S1), respectively.14 The strong
band at 1669 cm−1 and the medium one at 1605 cm−1 are attributed to the υ(C O) and
υ(C N) vibration.15 The band at 798 cm−1 is assigned to the υ(C S) vibration.15 The
S
2
H
3
N
1
N
C2H5
4
N
H
N
H
O
CH3
Figure 1 Acetoacetanilide-4-ethyl thiosemicarbazone.

367

368
R. R. ZAKY
O
PhHN
H3C
N
HN
NHC2H5
2H2O
S
H3C
O
O
Cu
CH3
O
O
OH2
Figure 2 Suggested structure of [Cu(H2L)(OAc)2(H2O)] · 2 H2O.
absence of bands at υ(OH) (∼ 3400 cm−1) and υ(SH) (2500–2600 cm−1) indicates that the
ligand exists in the keto and thione forms.16
Upon comparison of the spectrum of the ligand with that of [Cu(H2L)(OAc)2(H2O)] ·
2 H2O, it is clear that there is a negative shift of υ(C S) vibration (788 cm−1). On the other
hand υ(C O), υ(C N), υ(N1H), υ(N3H), and υ(N4H) vibrations remain more or less at
the same position (Table 2, Figure S1, available online in the Supplemental Materials).
This indicates that the coordination to the metal ion occurred through the sulfur atom of
the C S group, and the other groups do not participate in coordination. Thus, the ligand
acts as a neutral monodentate ligand via the thione sulfur atom (Figure 2).
In the case of the [Cu(H2L)Cl2(H2O)2] · H2O complex, H2L acts as a neutral bidentate
ligand coordinating via the carbonyl oxygen (C O) and azomethine nitrogen (C N)
atom. This mode of chelation is supported by the shift of both υ(C N) and υ(C O)
vibrations to lower wavenumber as shown in Table 1. The υ(C S), υ(N1H), υ(N3H), and
υ(N4H) vibrations are shifted to slightly higher wavenumbers, which indicates that they
do not participate in coordination to the metal ion. Also, new bands were observed at 510
and 405 cm−1 and assigned to υ(Cu O) and υ(Cu N),17 respectively. The spectrum of
the [Pd(H2L)Cl2] complex shows that the ligand behaves in the same manner as neutral
bidentate ligand (Table 2, Figure 3).
NHC2H5
X
HN
N
S
Cl
Cl
CH3
YH2O
M
O
X
NHPh
M = Cu(II), X = H2O,Y = 1
M = Pd(II), X = Y = 0
Figure 3 Suggested structure of [Cu(H2L)Cl2(H2O)2] · H2O and [Pd(H2L)Cl2].

369

370
R. R. ZAKY
NHC2H5
S
CH3
HN
N
O
S
H2O
O
Cu
PhHN
O
O
O
OH2
Figure 4 Suggested structure of [Cu(H2L)(SO4)(H2O)] . H2O.
On the other hand, the IR spectrum of the [Cu(H2L)(SO4)(H2O)] · H2O complex
shows a shift of υ(C N), υ(C O), and υ(C S) vibrations to lower wavenumbers. This
indicates that H2L acts as a neutral tridentate ligand coordinating via the carbonyl oxygen
(C O), the azomethine nitrogen (C N), and the thione sulfur (C S) atom. Also, the
spectrum of this complex shows new bands at 508 and 417 cm−1 assignable to υ(Cu O)
and υ(Cu N),17 respectively. The appearance of bands at 1058 and 920 cm−1 assignable
to the SO stretching vibrations suggests bidentate sulfate18 (Figure 4).
Moreover, the IR spectra indicate that H2L acts as a mononegative tridentate ligand
coordinating through the azomethine nitrogen (C N), the thione sulfur (C S), and the
deprotonated enolized carbonyl oxygen ( C O−) atom. This behavior is observed in
[Ni(HL)(OAc)(H2O)] · 2 H2O, [Co(HL)(OAc)(H2O)] · 2 H2O (Figure 5) and [UO2(HL)2] ·
H2O complexes. This mode of chelation is supported by: (i) the disappearance of υ(C O)
with simultaneous appearance of new bands in the 1110–1119 cm−1 region assignable
to υ(C O)enolic19; (ii) the shift of both υ(C N) and υ(C S) to lower wavenumber; (iii)
the υ(N1H), υ(N3H), and υ(N4H) vibrations are found more or less at same position,
which indicates that they do not participate in coordination to the metal ions; and (iv) the
appearance of new bands in the 506–515 and 405–421 cm−1 regions assignable to υ(M O)
and υ(M N). In addition, the spectrum of the [UO2(HL)2] · H2O complex displays two
NHC2H5
CH3
HN
S
N
O
2H2O
M
PhHN
CH3
O
O
OH2
M = Ni(II) and Co(II)
Figure 5 Suggested structure of [Ni(HL)(OAc)(H2O)] · 2 H2O and [Co(HL)(OAc)(H2O)] · 2 H2O.

ACETOACETANILIDE-4-ETHYL THIOSEMICARBAZONE COMPLEXES
371
bands at 906 and 840 cm−1 assigned to υ3 and υ1 vibrations of the dioxouranium ion,
respectively. The value of υ3 is used to calculate the force constant (F) of υ(U O) by the
method of McGlynn and Smith:20
(υ3)2 = (1307)2 (FU−O)/14.103
The force constant obtained for uranyl complex was then substituted into the relation given
by Jones21 to estimate the U O bond length in Å.
−1/3
RU−O = 1.08 (FU−O
)
+ 1.17
The calculated FU O and RU O values are 6.776 mdayn Å−1 and 1.740 Å, respectively,
and fall in the usual range for uranyl complexes.
Finally, the IR spectral data of [VO(L)(H2O)] · 2 H2O and [Cd(L)(H2O)] · H2O (Fig-
ures 6 and S2) show that H2L behaves as a binegative tridentate ligand coordinating to the
metal ions through the azomethine nitrogen atom, the deprotonated thiol sulfur atom, and
the deprotonated enolic carbonyl oxygen atom. This feature is supported by (i) the negative
shift of υ(C N) vibration; (ii) the disappearance of the υ(C O) vibration with simulta-
neous appearance of new bands in the 1130–1135 cm−1 range assignable to υ(C O)enolic
;
(iii) the disappearance of υ(C S) and υ(N3H) with simultaneous appearance of new bands
in the 610–620 and 1615–1625 cm−1 regions attributed to υ(C S) and υ(C N3),22 re-
spectively; (iv) the appearance of new bands in the 508–515 and 408–415 cm−1 regions
assignable to υ(M O) and υ(M N), respectively; and (v) a band at 976 cm−1 due to
υ(V O) vibration in agreement with the square-pyramidal arrangement around the metal
ion.23
In metal acetate complexes, the acetate group coordinates to the metal ions in a
bidentate manner, where the difference between the two acetate bands is ꢂν ≤ 180 cm−1 24
.
Also, the bands of coordinated water observed at 815–857 and 533–592 cm−1 are assigned
to ρr(H2O) and ρw(H2O), respectively.24 Moreover, strong evidence for the presence or
absence of water of crystallization and/or coordinated water is supported by the thermogram
of all complexes.
1
The H NMR spectroscopic data of H2L in d6-DMSO and of some complexes are
listed in Table 3. In the free ligand, the aromatic protons appear at δ 7.0–7.6 ppm, and the
signal at δ 2.0 ppm is assigned to the protons of the methylene group ( CH2 ), while the
signal of CH3 protons appears at δ 1.9 ppm. Also the two signals at δ 3.6 and 1.2 ppm are
ascribed to the protons of ethyl group. The N3H, N1H, and N4H protons appear at δ 10.6,
10.2, and 9.7 ppm, respectively.
The 1H NMR of the [Pd(H2L)Cl2] complex shows three signals at δ 10.5, 10.3, and
9.8 ppm assignable to the protons of N3H, N1H, and N4H, respectively. The chemical shifts
Table 3 1H NMR data of H2L and its diamagnetic complexes
Ethyl group
Aromatic
Compound
N3H N1H N4H
CH2
CH
CH3
protons
CH2
CH3
H2L
10.6 10.2
10.5 10.3
9.7
9.8
9.7
2.0
2.1
—
—
—
4.6
1,9
1.9
2.0
7.0–7.6
7.0–7.7
7.2–7.6
3.6
3.6
3.7
1.2
1.2
1.3
[Pd(H2L)Cl2]
[Cd(L)(H2O)] · H2O
—
10.3

372
R. R. ZAKY
Table 4 Magnetic moments, electronic spectra, and ligand field parameters of metal complexes of H2L
µeff
(BM)
Band position
(nm)
Dq
(cm−1
B
(cm−1
Compound
)
)
β
υ2/υ1
[Ni(HL)(OAc)(H2O)] · 2 H2O
[Co(HL)(OAc)(H2O)] · 2 H2O
[Cu(H2L)(OAc)2(H2O)] · 2 H2O
[Cu(H2L)Cl2(H2O)2] · H2O
[Cu(H2L)(SO4)(H2O)] · H2O
[VO(L)(H2O)] · 2 H2O
3.43
5.15
2.05
1.98
1.93
1.98
Diam.
Diam.
735; 440
711; 580
678; 600
682; 590
700; 600
850
903
763
—
—
—
—
—
—
813
803
—
—
—
—
—
—
0.78
0.83
—
—
—
—
—
—
1.67
2.15
—
—
—
—
—
—
[Pd(H2L)Cl2]
442
417; 359
[UO2(HL)2] · H2O
are almost the same as those of the free ligand, confirming that these groups are not involved
in coordination to the metal ion.
The 1H NMR spectrum of the [Cd(L)(H2O)] · H2O complex shows a new signal at δ
4.6, assigned to the proton of ( O C CH ) group, which is formed in the enolization of
the C O moiety. Also, the spectral data indicate the absence of a signal for the N3H proton
due to the thioenolization.
Electronic, ESR Spectra, and Magnetic Moment Measurements
The electronic spectrum of the ligand in DMF shows two bands at 305 and 350 nm
due to the π → π∗ and n → π∗ transitions, respectively.25 Previous studies proved that
the bands at 400–384 and 459–404 nm should be assigned to O → M26 and S → M27
transitions. The bands at 470–344 nm (DMF) in the spectra of the present complexes may
therefore be due to O →M and S → M CT.
The electronic spectrum of [VO(L)(H2O)] · 2 H2O shows a band at 580 nm (Table 4,
2
Figure S3), assigned to 2B2 → E (υ2) in a square-pyramidal configuration.28
The electronic spectrum of [Ni(HL)(OAc)(H2O)] · 2 H2O shows two bands at 735
3
3
3
3
and 440 nm attributed to A2g → T1g(F) and A2g → T1g(P) transitions (Figure 8),
respectively, in an octahedral geometry.29 The calculated values of Dq, B, β, and υ2/υ1
lie in the range reported for an octahedral structure.29 The position of υ1 (8146 cm−1
)
OH2
OH2
O
V
Cd
O
S
O
S
H2O
2H2O
NHPh
NHPh
N
C2H5HN
N
C2H5HN
N
N
H3C
H3C
Figure 6 Suggested structure of [Cd(L)(H2O)] · H2O and [VO(L)(H2O)] · 2 H2O complexes.

ACETOACETANILIDE-4-ETHYL THIOSEMICARBAZONE COMPLEXES
373
is calculated theoretically. The magnetic moment (µeff. = 3.43 BM) provides additional
evidence for an octahedral geometry.
The electronic spectrum of the [Co(HL)(OAc)(H2O)] · 2 H2O complex exhibits two
4
4
4
4
bands at 711 and 580 nm attributed to T1g → A2g(F) and T1g → T1g(P) transitions,
respectively, in an octahedral configuration.29 The calculated Dq, B, β, and υ2/υ1 values
are in the range reported for an octahedral environment around Co(II) complexes.[29] The
position of υ1 (6541cm−1) is calculated theoretically.29 Also, the value of the magnetic
moment (5.15 BM) is additional evidence for an octahedral geometry around the Co(II) ion.
The Cu(II) complexes have magnetic moment values (1.93–2.05 BM), which corre-
spond to those reported for d9-systems. The electronic spectra of [Cu(H2L)(OAc)2(H2O)] ·
2 H2O, [Cu(H2L)Cl2(H2O)2] · H2O and [Cu(H2L)(SO4)(H2O)] · H2O complexes show
broad bands in the 590–600 nm region with a shoulder in the 678–700 nm region, which
2
2
may be assigned to 2B1g → Eg and 2Eg → A1g transitions, respectively, in a tetragonally
distorted octahedral configuration.30
The electronic spectrum of the [UO2(HL)2] · H2O complex shows two bands at 417
1
2
and 359 nm, which may be ascribed to ꢃ+g→ πu transition for dioxouranium(VI) and
charge transfer n → π∗ transition, respectively.31
The diamagnetic [Pd(H2L)Cl2] complex32 shows a band at 442 nm, which is assigned
2
to 1A1g → B1g transition in a square-planar configuration.33
The room temperature solid state ESR spectra of copper complexes (Figure 7) exhibit
axially symmetric g-tensor parameters with g// > γ >2.0023. The g values reflect that the
⊥
Cu(II) center has a tetragonally distorted octahedral geometry with the unpaired electron
residing in the dx2-y2 orbital.34
No band corresponding to the forbidden magnetic dipolar transition for the complexes
was observed at half-field (ca. 1500 G, g = 4.0), ruling out any Cu–Cu interaction and
8000
c
6000
4000
2000
0
b
a
-2000
-4000
-6000
-8000
-10000
2500
3000
3500
4000
4500
G (Gauss)
Figure 7 Room temperature X-band ESR spectra of (a) [Cu(H2L)(OAc)2(H2O)] · 2 H2O, (b) [Cu(H2L)
Cl2(H2O)2] · H2O, and (c) [Cu(H2L)(SO4)(H2O)] · H2O.

374
R. R. ZAKY
Table 5 ESR data of the copper(II) and oxovanadium(IV) complexes at room temperature
A// × 10−4
Complex
g//
g
(cm−1
)
G
α2
β2
⊥
[Cu(H2L)(OAc)2(H2O)] · 2 H2O
[Cu(H2L)Cl2(H2O)2] · H2O
[Cu(H2L)(SO4)(H2O)] · H2O
[VO(L)(H2O)] · 2 H2O
2.27
2.26
2.23
1.95
2.04
2.06
2.05
1.97
155
160
160
195
5.6
4.3
4.6
—
0.75
0.76
0.75
0.73
0.82
0.75
0.65
0.92
indicating that the complex is a mononuclear Cu(II) complex. In axial symmetry, the
g-values are related by the expression, G = (g// − 2)/(γ –2) = 4, which measures the
⊥
exchange interaction between copper centers in the solid.
According to Hathaway and Billing,35 if the value of G is greater than 4, the exchange
interaction between the copper(II) centers in the solid state is negligible, whereas when
G is less than 4, a considerable exchange interaction is present in the solid complex. The
calculated G values are given in Table 5. The G values are greater than four, suggesting no
copper–copper exchange interaction, and these results are very consistent with the values
of the effective magnetic moments.
In hexacoordinated Cu(II) complexes, tetragonal distortion from the octahedral sym-
metry due to the Jahn–Teller distortion is very common. Such complexes with different
ligands are attractive mainly due to the variation in the coordination geometry and the
spectral features.36,37
Hyperfine splittings for the complexes were not observed at higher fields, excluding
any interaction of the nuclear spins of the nitrogen (I = 1) with the unpaired electron
density on Cu(II). The absence of 14N hyperfine coupling may be due to a relatively high
tetrahedral distortion for this complex.38
15000
10000
5000
0
-5000
-10000
2400 2600 2800 3000 3200 3400 3600 3800 4000 4200 4400 4600
G (Gauss)
Figure 8 Room temperature X-band ESR spectrum of [VO(L)(H2O)] · 2 H2O.

ACETOACETANILIDE-4-ETHYL THIOSEMICARBAZONE COMPLEXES
375
Molecular orbital coefficients, α2 (a measure for the covalency of the in-plane σ-
bonding between a copper 3D orbital and the ligand orbitals) and β2 (covalent in-plane
π-bonding), were calculated using the following equations,39–42 where α2 = 1 indicates
complete ionic character. On the other hand α2 = 0.5 denotes 100% covalent bonding with
the assumption of negligibly small values of the overlap integral:
ꢀ
ꢁ
A//
3(g − 2.0023)
⊥
α2 =
β2 =
+ (g// − 2.0023) +
+ 0.04
0.036
(g// − 2.0023)E
7
−8λα2
where λ = −828 cm−1 for the free copper ion and E is the electronic transition energy.
In Table 5, the α2 and β2 values indicate that there is a substantial interaction in the in-
plane σ-bonding, whereas the in-plane π-bonding is almost ionic. The higher value of β2
compared to that of α 2 indicates that the in-plane π-bonding is less covalent than the in-
plane σ-bonding. The α 2 value for the Cu(II) complex indicates a considerable covalency
in the bonding between the copper(II) ion and the ligand.
The room temperature (300 K) ESR spectrum of [VO(L)(H2O)] · 2 H2O gave a broad
signal with poorly resolved eight hyperfine lines (Figure 8) similar to those reported for
mononuclear vanadium molecule. In the powdered mono vanadium complex, the spectrum
showed parallel and perpendicular features, which indicates axially symmetric anisotropy
with well resolved 16-line hyperfine splitting characteristic for the interaction between the
electron and the vanadium nuclear spin (I = 7/2). The spin Hamiltonian parameters are
calculated and given in Table 5. The calculated ESR parameters indicate that the unpaired
electron (d1) in the complex is present in the dxy-orbital, in accord with square pyramidal
geometry.43 The values obtained for the complexes agree well with the g-tensor parameters
reported for distorted square pyramidal vanadium complexes.
The molecular orbital coefficients α2 and β2 for the complexes were calculated using
the following equations:44
ꢂ
ꢃ
4
3
A = −P k + β2 + (ge − g ) + (ge − g )
||
||
⊥
7
2
7
ꢂ
ꢃ
11
A
⊥
= −P k − β2 + (ge − g )
⊥
7
4
8α2β2λ
ge − g =
||
E
Neglecting the second order effects and taking negative values for A// and A , these
⊥
equations are solved to obtain α2 and β2.
The spin-orbit coupling coefficient λ is assumed to be 170 cm−1 for the M2+ ion, E
2
2
is the energy of B2 → E electronic transition, and k is the Fermi contact term, which is
directly related to the isotropic hyperfine coupling and represents the amount of unpaired
electron density at the nucleus. The lower values of α2 as compared to β2 indicate that the
in-plane σ-bonding is more covalent. These data are very consistent with other reported
data.45

376
R. R. ZAKY
Thermal Analysis
The TG curve of [VO(L)(H2O)] · 2 H2O complex displays 9.11% weight loss in the
temperature range 60–130◦C (Figure S4, Table 6), which correlates to two water molecules
outside the coordination sphere. In the temperature range 130–220◦C, the TG curve displays
4.55% weight loss, which could be correlated with the release of the coordinated water
molecule.46 Thermal degradation of the organic molecule starts above 220◦C. Over the
temperature range 220–462◦C, the weight loss of 26.57% in the TG curve is most probably
ascribed to the partial decomposition of the chelating agent with the elimination of the
loosely bonded (C2H6N + C3H3N2) moiety. The final weight loss of 28.09% ending at
653◦C is largely attributed to complete decomposition of the organic molecule alongside
rupture of the chelate bond, leaving VO2 + S and comprising 30.71% of the initial mass of
the complex.
The TG curve of the [Ni(HL)(OAc)(H2O)] · 2 H2O complex displays 8.14% weight
loss in the temperature range of 56–151◦C (Figure S5), which correlates to the two water
molecules outside of the coordination sphere. In the temperature range 151–245◦C, the
TG curve displays 17.34% weight loss, which could be correlated with the release of one
coordinated water molecule as well as of the acetate group.46,47 Thermal degradation of the
organic molecule starts above 245◦C. Over the temperature range of 245–439◦C, the weight
loss of 22.90% in the TG curve is most probably ascribed to the partial decomposition of
the chelating agent with the elimination of the loosely bonded C3H7N2S moiety. The final
weight loss of 35.13% ending at 650◦C is largely attributed to complete decomposition of
the organic molecule alongside rupture of the chelate bond, leaving NiO and comprising
16.60% of the initial mass of the complex.
For more information about the effects of the anions on the structural properties
and bonding, the coefficients of the thermal behavior of the complexes and the heat of
activation E# of the decomposition stages were determined from the TG thermograms for
some steps of the decomposition of the VO2+ and Ni2+ complexes using the Coats–Redfern
equations.48 The other kinetic parameters, ꢂH#, ꢂS#, and ꢂG#, were computed using the
relationships: ꢂH# = E#−RT, ꢂS# = R[ln(Ah/kT)−1], and ꢂG# = ꢂH# −TꢂS#, where
k is Boltzmann’s constant and h is Planck’s constant. The activation parameters are listed
in Table 6. It has been found that some of the steps are unsuitable for kinetic analysis.
The overlapping steps reflect slight differences between the calculated and found weight
losses.
Some of the activation entropies ꢂS# are negative, indicating that the reactions are
slow or the activated complex is more ordered than the reactants.49 The positive values of
ꢂH# mean that the decomposition processes are endothermic. All decomposition stages
of the complexes show a best fit for (n = 1), indicating a first-order decomposition in all
cases. Other n values (e.g., 0, 0.33, and 0.66) did not lead to better correlations. The values
of ꢂG# increase significantly for the subsequent decomposition stages of a given complex
due to increasing of TꢂS#.
Antimicrobial Activity Bioassay
The antimicrobial activity of the investigated ligand and complexes against Bacillus
Subtilis and Micrococcus sp. is summarized in Table S1 (Supplemental Materials)

377

378
R. R. ZAKY
CONCLUSION
The structures of the complexes of H2L with Co2+, Ni2+, Cu2+, Pd2+, Cd2+, UO2
,
2+
and VO2+ ions were determined by elemental analyses, IR, 1H NMR spectroscopy, molar
conductance, magnetic, UV-Vis, ESR, and thermal analysis data. From the IR spectra, it
is concluded that H2L behaves as a monodentate, bidentate, and/or tridentate ligand. From
the molar conductance data, it was found that all the complexes are non-electrolytes. The
1H NMR spectrum of the free ligand shows that the OH and SH signals are absent. This
indicates that the ligand exists in the keto and thione forms. On the basis of magnetic
measurements, UV-Vis, and ESR, an octahedral geometry is suggested for all investigated
complexes, except for the Pd(II) complex, which is square planar, and for the VO2+
complex, which is square-pyramidal. The ligand and the metal complexes prepared were
also screened for their antibacterial activity against bacterial species Bacillus subtilis as
Gram positive and Micrococcus sp as Gram negative bacteria. The activity data show that
most of the metal complexes were more potent antibactrials than the parent ligand against
these bacterial species. Moreover, the degradation powers of the tested compounds on the
calf thymus DNA were high.
EXPERIMENTAL
Apparatus and Reagents
The IR absorption spectra were recorded with a Mattson 5000 FTIR spectropho-
tometer. The electronic spectra were measured with a Unicam UV/VIS spectrometer UV2.
Thermogravimetric analysis was performed using an automatic recording thermobalance
type 951 (DuPont instrument). Samples were subjected to heat at a rate of 10◦C/min
(25–800◦C) in N2. 1H NMR spectra for H2L and diamagnetic complexes in DMSO were
recorded with an EM-390 (200 MHz) spectrometer. Carbon and hydrogen content for the
ligand and its complexes was determined at the Microanalytical Unit, Mansoura University,
Egypt. Co2+, Ni2+, Cu2+, Pd2+, Cd2+, UO22+, VO2+, Cl, and SO4 contents in the com-
2
plexes were determined by well known standard methods.50 ESR spectra were obtained
with a Bruker EMX spectrometer working in the X-band (9.78 GHz) with 100 kHz mod-
ulation frequency. The microwave power was set at 1 mW, and modulation amplitude was
set at 4 Gauss. The low field signal was obtained after 4 scans with a 10-fold increase in
the receiver gain. A powder spectrum was obtained in a 2 mm quartz capillary at room
temperature. All metal salts used were pure (Fluka, Aldrich, or Merck).
Preparation of the Ligand
Acetoacetanilide-4-ethyl thiosemicarbazone (Figure 1) was prepared by stirring a
mixture of 4-ethyl-3-thiosemicarbazide (1.19 g, 0.1 mol) with acetoacetanilide (1.77 g, 0.1
mol) in EtOH (40 mL) and adding 2 mL of 5M HCl. The reaction mixture was stirred at
room temperature for 4 h. The white precipitate was filtered off, washed with ethanol and
Et2O, and finally recrystallized from ethanol (mp 110◦C; yield 90%). The purity of the
compound was checked by TLC.

ACETOACETANILIDE-4-ETHYL THIOSEMICARBAZONE COMPLEXES
379
Preparation of the Complexes
The complexes were prepared by mixing equimolar amounts of H2L with ethanolic
and/or aqueous solution of the chloride salt of Cu2+, Cd2+, and Pd2+ as potassium tetra-
chloropalladate; of the acetate salt of Ni2+, Co2+, Cu2+, and UO22+; and the sulfate salt of
Cu2+ and VO2+. In the case of VO2+ and Cd2+, 0.1 g of sodium acetate was added as a
buffering agent to raise the pH medium. The reaction mixture was heated under reflux on a
water bath for 1–3 h. The precipitate was filtered off, washed with hot EtOH and/or H2O,
and finally preserved in a vacuum desiccator over anhydrous CaCl2.
Biological Activity (See Supplemental Materials)
Antimicrobial activity. The ligand and its complexes have been tested as antibac-
terial agents at the Department of Microbiology, Faculty of Pharmacy, Mansoura University.
Inhibition zone was measured51 and compared with that of gentamicin solution (commercial
antibiotic, Memphis Co., Egypt, 1000 µg mL−1). The experimental control was DMF.
Genotoxicity activity. A solution of 2 mg of calf thymus DNA was dissolved in
1 mL of sterile distilled water (see Figure S6).
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