# Synthesis, Characterization, and Anticancer Activity of  New Metal Complexes Derived from 2-Hydroxy-3-(hydroxyimino)-4-oxopentan-2-ylidene)benzohydrazide

Authors: Abdou Saad El-Tabl; Moshira Mohamed Abd El-Waheed; Mohammed Ahmed Wahba; Nahla Abd El-Halim Abou El-Fadl
Journal: Bioinorganic Chemistry and Applications
 (2015)

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## Abstract

Novel metal(II) complexes derived from 2-hydroxy-N′-((Z)-3-(hydroxyimino)-4-oxopentan-2-ylidene)benzohydrazide ligand (H2L) were synthesized and characterized by elemental and thermal analyses (DTA and TGA), IR, UV-VIS, 1H-NMR, ESR and mass spectroscopy, magnetic susceptibilities, and conductivities measurements. The complexes adopt distorted octahedral geometry. The ESR spectra of the solid copper(II) complexes are characteristic to d9 configuration and have an axial symmetry type of a d(x2-y2) ground state. The g values confirmed the tetragonal octahedral geometry with a considerably ionic or covalent environment. The cytotoxic activity of the ligand and its metal complexes showed potent cytotoxicity effect against growth of human liver cancer HepG2 cell lines compared to the clinically used Sorafenib (Nexavar).


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## Body

## 1. Introduction

There is a growing interest in oxime-hydrazone and their coordination compounds caused by their biological activity [1, 2]. Many clinically successful anticancer drugs were either naturally occurring molecules or have been developed from their synthetic analogs. Metal complexes have unique properties enhancing their role as antitumor agents. An important property is the ability of metals to form positively charged ions in an aqueous solution that can bind to negatively charged biological molecules [3, 4]. The high electron affinity of metal ions can significantly polarize groups that are coordinated to them, leading to the generation of hydrolysis reactions [4]. Furthermore, metal ions also has the ability to coordinate ligands in a three dimensional configuration, thus allowing functionalization of groups that can be tailored to defined molecular targets [5, 6]. Recently, considerable attention has been drawn to oximes, hydrazones, and their coordinated compounds due to their biological activities as fungicides [7, 8], bactericides [9], analgesic and anti-inflammatory [10], antioxidant [11, 12], antitumor [13–15], and insecticidal [16]. Metal complexes of bis-hydrazone derived from isatin monohydrazone and 2-hydroxy-l-naphthaldehyde have been reported and they demonstrated interesting biological properties [17]. Oxovanadium(IV) complexes derived from 2-thiophene carboxylic acid hydrazide showed a great cytotoxicity towards Artemia salina [18]. Homo- and heteronuclear copper(II) and nickel(II) complexes derived from oxime-type ligands have been also reported; the observed IC50 values indicated that they are potential antioxidant [19]. Cytotoxicity of a series of cobalt(II) complexes of 2-furaldehyde oximes was compared with copper complexes of furan oximes to determine whether the type of metal is important to the cytotoxicity and mode of action of the complexes. It was shown that varying the type of metal produces differences in both cytotoxicity and mode of action [20]. Oxime hydrazones can react with metal(II) salts to produce either mono- or binuclear complexes. The keto hydrazone moiety may coordinate to metals in the ketoamide or deprotonated enolimine form. Compounds containing both oxime and hydrazone groups are typically act as tridentate, mono- or biprotic ligands coordinating through the amide oxygen, imine, and oxime nitrogen atom depending on the reaction conditions [21]. Considerable interest has been attracted to synthesize oxime-hydrazide compounds as important target structures and evaluated their biological activities. These observations have been guiding the development of new compounds that possess varied biological activities. In view of interest and importance of oxime-hydrazide complexes, we reported here synthesis and characterization of new metal complexes derived from 2-hydroxy-N′-((Z)-3-(hydroxyimino)-4-oxopentan-2-ylidene)benzohydrazide. The work was extended to study the anticancer activity of the ligand and its metal complexes against human liver cancer HepG2.


## 2. Experimental

### 2.1. Instrumentation

All reagents employed for the preparation of the ligand and its complexes were of the analytical grade available and used without further purification. Metal salts and salicylic hydrazide were provided from SIGMA-ALDRICH company; diacetylmonoxime was prepared by a published method [22]. The purity of all compounds was confirmed by TLC. The ligand and its metal complexes were analyzed for C, H, and N at the Microanalytical center, Cairo University, Egypt. Standard analytical methods were used to determine the metal ion content [23]. FT-IR spectra of the ligand and its metal complexes were measured using KBr discs by a Jasco FT/IR 300E Fourier transform infrared spectrophotometer covering the range 400–4000 cm−1. Electronic spectra in the 200–900 nm regions were recorded on a Perkin-Elmer 550 spectrophotometer. The thermal analyses (DTA and TGA) were carried out on a Shimadzu DT-30 thermal analyzer from room temperature to 800°C at a heating rate of 10°C/min. Magnetic susceptibilities were measured at 25°C by the Gouy method using mercuric tetrathiocyanatocobaltate(II) as the magnetic susceptibility standard. Diamagnetic corrections were estimated from Pascal’s constant [24]. The magnetic moments were calculated from the equation(1)μeff.=2.84χMcorr·T.The molar conductance of 10−3 M solution of the complexes in DMSO was measured at 25°C with a Bibby conductometer type MCl. The resistance measured in ohms and the molar conductivities were calculated according to the equation(2)ΛM=V∗K∗gMw∗Ω,where ΛM = molar conductivity (Ω−1 cm2 mol−1), V = volume of the complex solution (mL), K = cell constant (0.92/cm−1), Mw = molecular weight of the complex, g = weight of the complex (g), Ω = resistance (Ω). 1H-NMR spectra of the ligand, and its Zn(II), Cd(II), and Hg(II) complexes were obtained on Perkin-Elmer R32-90-MHz spectrophotometer. Chemical shifts (ppm) were reported relative to TMS. ESR measurements of solid complexes at room temperature were made using a Varian E-109 spectrophotometer, using DPPH as a standard material. Mass spectra were recorded using JEULJMS-AX-500 mass spectrometer.


### 2.2. Preparation of the Ligand and Its Metal Complexes

#### 2.2.1. Preparation of the Ligand [H2L](1)

The ligand (H2L) was prepared by dropwise addition of equimolar amounts of salicylic hydrazide (2-hydroxy-benzohydrazide) (1.52 g, 0.01 mol) dissolved in 20 mL of absolute ethanol to an ethanolic solution of diacetyl monoxime (1.29 g, 0.01 mol) (Figure 1). The mixture was refluxed with stirring for 4 hrs. A dark green precipitate was obtained, filtered off, washed with ethanol, and dried under vacuum over P2O5. Analytical data of the ligand are given in (Table 1).Table 1
Analytical and physical data of the ligand [H2L] (1) and its metal complexes.
Comp. number
Molecular formula
Color
FW
M.P (°C)
Yield (%)
Anal./found (calc.) (%)
Molar conductance∗
C
H
N
M
Cl
(1)
[H2L]   C12H13N3O4
Dark green
263
120
80
55.01 (54.75)
5.48 (4.98)
16.01 (15.96)
—
—
(2)
[(HL)Cu(OAc)(H2O)2]·H2O   C14H21N3O9Cu
Green
439
>300
85
39.02 (38.31)
4.61 (4.82)
9.33 (9.57)
14.50 (14.48)
—
10.12
(3)
[(HL)Cu2(OAc)3(H2O)5]·H2O  C18H33N3O15Cu2
Greenish brown
658.56
>300
90
33.00 (32.83)
5.11 (5.05)
6.51 (6.38)
19.81 (19.30)
—
13.15
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
Yellowish green
762
>300
90
44.47 (44.12)
4.39 (5.03)
11.22 (11.03)
8.52 (8.34)
—
11.30
(5)
[(H2L)CuCl2(H2O)]·2H2O   C12H19Cl2N3O7Cu
Brownish green
452
>300
85
32.10 (31.90)
4.23 (4.24)
9.52 (9.30)
14.30 (14.07)
16.00 (15.70)
14.10
(6)
[(H2L)Cu(SO4)(H2O)2]   C12H17N3O10SCu
Dark olive
459
>300
80
31.71 (31.41)
4.00 (3.73)
9.72 (9.16)
14.10 (13.85)
—
16.20
(7)
[(HL)Ni(OAc)(H2O)2]·2H2O   C14H23N3O10Ni
Pale brown
452
>300
70
37.59 (37.20)
4.71 (5.13)
9.82 (9.30)
13.21 (12.98)
—
10.87
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
Yellow
487
>300
90
32.32 (32.06)
4.81 (4.55)
9.10 (8.63)
12.52 (12.05)
—
15.70
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
Dark brown
470
>300
78
35.98 (35.75)
5.51 (5.36)
9.12 (8.93)
13.00 (12.53)
—
11.81
(10)
[(H2L)Co(SO4)(H2O)2]·3H2O   C13H26N3O13SCo
Pale white
523
>300
70
30.12 (29.83)
5.31 (5.01)
8.52 (8.03)
11.82 (11.26)
—
15.87
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
Brown
466
>300
75
35.66 (36.06)
5.12 (5.40)
9.33 (9.01)
12.12 (11.78)
—
11.12
(12)
[(HL)Zn(OAc)(H2O)2]·2H2O   C14H23N3O10Zn
Brown green
459
>300
75
37.00 (36.65)
5.51 (5.05)
9.72 (9.16)
14.73 (14.25)
—
11.87
(13)
[(H2L)Zn(SO4)(H2O)2]·3H2O   C12H23N3O13SZn
Reddish brown
515
>300
70
28.33 (28.01)
4.62 (4.500)
8.51 (8.16)
13.11 (12.70)
—
15.72
(14)
[(HL)Cd(OAc)(H2O)2]·2H2O   C14H23N3O10Cd
White
506
>300
80
33.72 (33.25)
4.72 (4.58)
8.71 (8.31)
22.87 (22.23)
—
12.11
(15)
[(H2L)Cd(SO4)(H2O)2]·2H2O   C12H21N3O12SCd
Yellowish white
544
>300
85
27.01 (26.50)
3.95 (3.89)
8.11 (7.73)
21.10 (20.67)
—
17.11
(16)
[(HL)Hg(OAc)]·H2O   C14H17N3O7Hg
Yellowish green
540
>300
70
31.52 (31.15)
3.32 (3.17)
8.11 (7.78)
37.52 (37.15)
—
9.82
(17)
[(HL)Fe(Cl)2(H2O)]·3H2O   C12H20N3O8Cl2Fe
Black
461
>300
75
31.65 (31.26)
4.72 (4.37)
9.35 (9.11)
12.45 (12.11)
15.87 (15.38)
14.87
∗
Λm (Ω−1 cm2 mol−1).Figure 1
Preparation of the ligand [H2L].


#### 2.2.2. Preparation of Metal Complexes,(2)–(17)

Complexes(2)–(17) were synthesized by refluxing 25 mL ethanoic solution of the ligand with 25 mL ethanolic solution of 3.79 g, 0.02 mol of Cu(OAc)2·H2O (1 L : 1 M), complex(2); 7.58 g, 0.04 mol of Cu(OAc)2·H2O (1 L : 2 M), complex(3); 1.89 g, 0.01 mol of Cu(OAc)2·H2O (2 L : 1 M), complex(4); 1.88 g, 0.02 mol of CuCl2·2H2O (1 L : 1 M), complex(5); 3.03 g, 0.02 mol of CuSO4·5H2O (1 L : 1 M), complex(6); 4.73 g, 0.02 mol of Ni(OAc)2·4H2O (1 L : 1 M), complex(7); 4.99 g, 0.02 mol of NiSO4·6H2O (1 L : 1 M), complex(8); 4.73 g, 0.02 mol of Co(OAc)2·4H2O (1 L : 1 M), complex(9); 2.94 g, 0.02 mol of CoSO4·4H2O (1 L : 1 M), complex(10); 5.86 g, 0.02 mol of 4.65 g, 0.02 mol of Mn(OAc)2·4H2O (1 L : 1 M), complex(11); 4.17 g, 0.02 mol of Zn(OAc)2·2H2O (1 L : 1 M), complex(12); 3.41 g, 0.02 mol of ZnSO4·4H2O (1 L : 1 M), complex(13); 5.06 g, 0.02 mol of Cd(OAc)2·2H2O (1 L : 1 M), complex(14); 4.87 g, 0.02 mol of CdSO4·8H2O (1 L : 1 M), complex(15); 6.05 g, 0.02 mol of Hg(OAc)2·4H2O (1 L : 1 M), complex(16); 5.13 g, 0.02 mol of FeCl3·6H2O (1 L : 1 M), complex(17). The reaction mixtures were refluxed with stirring for 2–4 hrs range, depending on the nature of the metal ion and the anion. The precipitates so formed were filtrated off, washed with ethanol, and dried in vacuum over P2O5. Analytical data of the metal complexes are given in (Table 1).


### 2.3. Biological Activity

Evaluation of the cytotoxic activity of the ligand and its metal complexes was carried out in the Pathology Department, Faculty of Medicine, El-Menoufia University, Egypt. The evaluation process was carried out in vitro using the Sulfo-Rhodamine-B-stain (SRB) [25]. Cells were plated in 96-multiwell plate (104 cells/well) for 24 hrs before treatment with the complexes to allow attachment of cell to the wall of the plate. Different concentrations of the compounds under test in DMSO (0, 5, 12.5, 25, and 50 μg/mL) were added to the cell monolayer, triplicate wells being prepared for each individual dose. Monolayer cells were incubated with the complexes for 48 hrs at 37°C and in atmosphere of 5% CO2. After 48 hrs, cells were fixed, washed, and stained with Sulfo-Rhodamine-B-stain. Excess stain was wash with acetic acid and attached stain was recovered with Tris EDTA buffer. Color intensity was measured in an ELISA reader. The relation between surviving fraction and drug concentration is plotted to get the survival curve for each tumor cell line after addition the specified compound.


## 2.1. Instrumentation

All reagents employed for the preparation of the ligand and its complexes were of the analytical grade available and used without further purification. Metal salts and salicylic hydrazide were provided from SIGMA-ALDRICH company; diacetylmonoxime was prepared by a published method [22]. The purity of all compounds was confirmed by TLC. The ligand and its metal complexes were analyzed for C, H, and N at the Microanalytical center, Cairo University, Egypt. Standard analytical methods were used to determine the metal ion content [23]. FT-IR spectra of the ligand and its metal complexes were measured using KBr discs by a Jasco FT/IR 300E Fourier transform infrared spectrophotometer covering the range 400–4000 cm−1. Electronic spectra in the 200–900 nm regions were recorded on a Perkin-Elmer 550 spectrophotometer. The thermal analyses (DTA and TGA) were carried out on a Shimadzu DT-30 thermal analyzer from room temperature to 800°C at a heating rate of 10°C/min. Magnetic susceptibilities were measured at 25°C by the Gouy method using mercuric tetrathiocyanatocobaltate(II) as the magnetic susceptibility standard. Diamagnetic corrections were estimated from Pascal’s constant [24]. The magnetic moments were calculated from the equation(1)μeff.=2.84χMcorr·T.The molar conductance of 10−3 M solution of the complexes in DMSO was measured at 25°C with a Bibby conductometer type MCl. The resistance measured in ohms and the molar conductivities were calculated according to the equation(2)ΛM=V∗K∗gMw∗Ω,where ΛM = molar conductivity (Ω−1 cm2 mol−1), V = volume of the complex solution (mL), K = cell constant (0.92/cm−1), Mw = molecular weight of the complex, g = weight of the complex (g), Ω = resistance (Ω). 1H-NMR spectra of the ligand, and its Zn(II), Cd(II), and Hg(II) complexes were obtained on Perkin-Elmer R32-90-MHz spectrophotometer. Chemical shifts (ppm) were reported relative to TMS. ESR measurements of solid complexes at room temperature were made using a Varian E-109 spectrophotometer, using DPPH as a standard material. Mass spectra were recorded using JEULJMS-AX-500 mass spectrometer.


## 2.2. Preparation of the Ligand and Its Metal Complexes

### 2.2.1. Preparation of the Ligand [H2L](1)

The ligand (H2L) was prepared by dropwise addition of equimolar amounts of salicylic hydrazide (2-hydroxy-benzohydrazide) (1.52 g, 0.01 mol) dissolved in 20 mL of absolute ethanol to an ethanolic solution of diacetyl monoxime (1.29 g, 0.01 mol) (Figure 1). The mixture was refluxed with stirring for 4 hrs. A dark green precipitate was obtained, filtered off, washed with ethanol, and dried under vacuum over P2O5. Analytical data of the ligand are given in (Table 1).Table 1
Analytical and physical data of the ligand [H2L] (1) and its metal complexes.
Comp. number
Molecular formula
Color
FW
M.P (°C)
Yield (%)
Anal./found (calc.) (%)
Molar conductance∗
C
H
N
M
Cl
(1)
[H2L]   C12H13N3O4
Dark green
263
120
80
55.01 (54.75)
5.48 (4.98)
16.01 (15.96)
—
—
(2)
[(HL)Cu(OAc)(H2O)2]·H2O   C14H21N3O9Cu
Green
439
>300
85
39.02 (38.31)
4.61 (4.82)
9.33 (9.57)
14.50 (14.48)
—
10.12
(3)
[(HL)Cu2(OAc)3(H2O)5]·H2O  C18H33N3O15Cu2
Greenish brown
658.56
>300
90
33.00 (32.83)
5.11 (5.05)
6.51 (6.38)
19.81 (19.30)
—
13.15
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
Yellowish green
762
>300
90
44.47 (44.12)
4.39 (5.03)
11.22 (11.03)
8.52 (8.34)
—
11.30
(5)
[(H2L)CuCl2(H2O)]·2H2O   C12H19Cl2N3O7Cu
Brownish green
452
>300
85
32.10 (31.90)
4.23 (4.24)
9.52 (9.30)
14.30 (14.07)
16.00 (15.70)
14.10
(6)
[(H2L)Cu(SO4)(H2O)2]   C12H17N3O10SCu
Dark olive
459
>300
80
31.71 (31.41)
4.00 (3.73)
9.72 (9.16)
14.10 (13.85)
—
16.20
(7)
[(HL)Ni(OAc)(H2O)2]·2H2O   C14H23N3O10Ni
Pale brown
452
>300
70
37.59 (37.20)
4.71 (5.13)
9.82 (9.30)
13.21 (12.98)
—
10.87
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
Yellow
487
>300
90
32.32 (32.06)
4.81 (4.55)
9.10 (8.63)
12.52 (12.05)
—
15.70
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
Dark brown
470
>300
78
35.98 (35.75)
5.51 (5.36)
9.12 (8.93)
13.00 (12.53)
—
11.81
(10)
[(H2L)Co(SO4)(H2O)2]·3H2O   C13H26N3O13SCo
Pale white
523
>300
70
30.12 (29.83)
5.31 (5.01)
8.52 (8.03)
11.82 (11.26)
—
15.87
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
Brown
466
>300
75
35.66 (36.06)
5.12 (5.40)
9.33 (9.01)
12.12 (11.78)
—
11.12
(12)
[(HL)Zn(OAc)(H2O)2]·2H2O   C14H23N3O10Zn
Brown green
459
>300
75
37.00 (36.65)
5.51 (5.05)
9.72 (9.16)
14.73 (14.25)
—
11.87
(13)
[(H2L)Zn(SO4)(H2O)2]·3H2O   C12H23N3O13SZn
Reddish brown
515
>300
70
28.33 (28.01)
4.62 (4.500)
8.51 (8.16)
13.11 (12.70)
—
15.72
(14)
[(HL)Cd(OAc)(H2O)2]·2H2O   C14H23N3O10Cd
White
506
>300
80
33.72 (33.25)
4.72 (4.58)
8.71 (8.31)
22.87 (22.23)
—
12.11
(15)
[(H2L)Cd(SO4)(H2O)2]·2H2O   C12H21N3O12SCd
Yellowish white
544
>300
85
27.01 (26.50)
3.95 (3.89)
8.11 (7.73)
21.10 (20.67)
—
17.11
(16)
[(HL)Hg(OAc)]·H2O   C14H17N3O7Hg
Yellowish green
540
>300
70
31.52 (31.15)
3.32 (3.17)
8.11 (7.78)
37.52 (37.15)
—
9.82
(17)
[(HL)Fe(Cl)2(H2O)]·3H2O   C12H20N3O8Cl2Fe
Black
461
>300
75
31.65 (31.26)
4.72 (4.37)
9.35 (9.11)
12.45 (12.11)
15.87 (15.38)
14.87
∗
Λm (Ω−1 cm2 mol−1).Figure 1
Preparation of the ligand [H2L].


### 2.2.2. Preparation of Metal Complexes,(2)–(17)

Complexes(2)–(17) were synthesized by refluxing 25 mL ethanoic solution of the ligand with 25 mL ethanolic solution of 3.79 g, 0.02 mol of Cu(OAc)2·H2O (1 L : 1 M), complex(2); 7.58 g, 0.04 mol of Cu(OAc)2·H2O (1 L : 2 M), complex(3); 1.89 g, 0.01 mol of Cu(OAc)2·H2O (2 L : 1 M), complex(4); 1.88 g, 0.02 mol of CuCl2·2H2O (1 L : 1 M), complex(5); 3.03 g, 0.02 mol of CuSO4·5H2O (1 L : 1 M), complex(6); 4.73 g, 0.02 mol of Ni(OAc)2·4H2O (1 L : 1 M), complex(7); 4.99 g, 0.02 mol of NiSO4·6H2O (1 L : 1 M), complex(8); 4.73 g, 0.02 mol of Co(OAc)2·4H2O (1 L : 1 M), complex(9); 2.94 g, 0.02 mol of CoSO4·4H2O (1 L : 1 M), complex(10); 5.86 g, 0.02 mol of 4.65 g, 0.02 mol of Mn(OAc)2·4H2O (1 L : 1 M), complex(11); 4.17 g, 0.02 mol of Zn(OAc)2·2H2O (1 L : 1 M), complex(12); 3.41 g, 0.02 mol of ZnSO4·4H2O (1 L : 1 M), complex(13); 5.06 g, 0.02 mol of Cd(OAc)2·2H2O (1 L : 1 M), complex(14); 4.87 g, 0.02 mol of CdSO4·8H2O (1 L : 1 M), complex(15); 6.05 g, 0.02 mol of Hg(OAc)2·4H2O (1 L : 1 M), complex(16); 5.13 g, 0.02 mol of FeCl3·6H2O (1 L : 1 M), complex(17). The reaction mixtures were refluxed with stirring for 2–4 hrs range, depending on the nature of the metal ion and the anion. The precipitates so formed were filtrated off, washed with ethanol, and dried in vacuum over P2O5. Analytical data of the metal complexes are given in (Table 1).


## 2.2.1. Preparation of the Ligand [H2L](1)

The ligand (H2L) was prepared by dropwise addition of equimolar amounts of salicylic hydrazide (2-hydroxy-benzohydrazide) (1.52 g, 0.01 mol) dissolved in 20 mL of absolute ethanol to an ethanolic solution of diacetyl monoxime (1.29 g, 0.01 mol) (Figure 1). The mixture was refluxed with stirring for 4 hrs. A dark green precipitate was obtained, filtered off, washed with ethanol, and dried under vacuum over P2O5. Analytical data of the ligand are given in (Table 1).Table 1
Analytical and physical data of the ligand [H2L] (1) and its metal complexes.
Comp. number
Molecular formula
Color
FW
M.P (°C)
Yield (%)
Anal./found (calc.) (%)
Molar conductance∗
C
H
N
M
Cl
(1)
[H2L]   C12H13N3O4
Dark green
263
120
80
55.01 (54.75)
5.48 (4.98)
16.01 (15.96)
—
—
(2)
[(HL)Cu(OAc)(H2O)2]·H2O   C14H21N3O9Cu
Green
439
>300
85
39.02 (38.31)
4.61 (4.82)
9.33 (9.57)
14.50 (14.48)
—
10.12
(3)
[(HL)Cu2(OAc)3(H2O)5]·H2O  C18H33N3O15Cu2
Greenish brown
658.56
>300
90
33.00 (32.83)
5.11 (5.05)
6.51 (6.38)
19.81 (19.30)
—
13.15
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
Yellowish green
762
>300
90
44.47 (44.12)
4.39 (5.03)
11.22 (11.03)
8.52 (8.34)
—
11.30
(5)
[(H2L)CuCl2(H2O)]·2H2O   C12H19Cl2N3O7Cu
Brownish green
452
>300
85
32.10 (31.90)
4.23 (4.24)
9.52 (9.30)
14.30 (14.07)
16.00 (15.70)
14.10
(6)
[(H2L)Cu(SO4)(H2O)2]   C12H17N3O10SCu
Dark olive
459
>300
80
31.71 (31.41)
4.00 (3.73)
9.72 (9.16)
14.10 (13.85)
—
16.20
(7)
[(HL)Ni(OAc)(H2O)2]·2H2O   C14H23N3O10Ni
Pale brown
452
>300
70
37.59 (37.20)
4.71 (5.13)
9.82 (9.30)
13.21 (12.98)
—
10.87
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
Yellow
487
>300
90
32.32 (32.06)
4.81 (4.55)
9.10 (8.63)
12.52 (12.05)
—
15.70
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
Dark brown
470
>300
78
35.98 (35.75)
5.51 (5.36)
9.12 (8.93)
13.00 (12.53)
—
11.81
(10)
[(H2L)Co(SO4)(H2O)2]·3H2O   C13H26N3O13SCo
Pale white
523
>300
70
30.12 (29.83)
5.31 (5.01)
8.52 (8.03)
11.82 (11.26)
—
15.87
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
Brown
466
>300
75
35.66 (36.06)
5.12 (5.40)
9.33 (9.01)
12.12 (11.78)
—
11.12
(12)
[(HL)Zn(OAc)(H2O)2]·2H2O   C14H23N3O10Zn
Brown green
459
>300
75
37.00 (36.65)
5.51 (5.05)
9.72 (9.16)
14.73 (14.25)
—
11.87
(13)
[(H2L)Zn(SO4)(H2O)2]·3H2O   C12H23N3O13SZn
Reddish brown
515
>300
70
28.33 (28.01)
4.62 (4.500)
8.51 (8.16)
13.11 (12.70)
—
15.72
(14)
[(HL)Cd(OAc)(H2O)2]·2H2O   C14H23N3O10Cd
White
506
>300
80
33.72 (33.25)
4.72 (4.58)
8.71 (8.31)
22.87 (22.23)
—
12.11
(15)
[(H2L)Cd(SO4)(H2O)2]·2H2O   C12H21N3O12SCd
Yellowish white
544
>300
85
27.01 (26.50)
3.95 (3.89)
8.11 (7.73)
21.10 (20.67)
—
17.11
(16)
[(HL)Hg(OAc)]·H2O   C14H17N3O7Hg
Yellowish green
540
>300
70
31.52 (31.15)
3.32 (3.17)
8.11 (7.78)
37.52 (37.15)
—
9.82
(17)
[(HL)Fe(Cl)2(H2O)]·3H2O   C12H20N3O8Cl2Fe
Black
461
>300
75
31.65 (31.26)
4.72 (4.37)
9.35 (9.11)
12.45 (12.11)
15.87 (15.38)
14.87
∗
Λm (Ω−1 cm2 mol−1).Figure 1
Preparation of the ligand [H2L].


## 2.2.2. Preparation of Metal Complexes,(2)–(17)

Complexes(2)–(17) were synthesized by refluxing 25 mL ethanoic solution of the ligand with 25 mL ethanolic solution of 3.79 g, 0.02 mol of Cu(OAc)2·H2O (1 L : 1 M), complex(2); 7.58 g, 0.04 mol of Cu(OAc)2·H2O (1 L : 2 M), complex(3); 1.89 g, 0.01 mol of Cu(OAc)2·H2O (2 L : 1 M), complex(4); 1.88 g, 0.02 mol of CuCl2·2H2O (1 L : 1 M), complex(5); 3.03 g, 0.02 mol of CuSO4·5H2O (1 L : 1 M), complex(6); 4.73 g, 0.02 mol of Ni(OAc)2·4H2O (1 L : 1 M), complex(7); 4.99 g, 0.02 mol of NiSO4·6H2O (1 L : 1 M), complex(8); 4.73 g, 0.02 mol of Co(OAc)2·4H2O (1 L : 1 M), complex(9); 2.94 g, 0.02 mol of CoSO4·4H2O (1 L : 1 M), complex(10); 5.86 g, 0.02 mol of 4.65 g, 0.02 mol of Mn(OAc)2·4H2O (1 L : 1 M), complex(11); 4.17 g, 0.02 mol of Zn(OAc)2·2H2O (1 L : 1 M), complex(12); 3.41 g, 0.02 mol of ZnSO4·4H2O (1 L : 1 M), complex(13); 5.06 g, 0.02 mol of Cd(OAc)2·2H2O (1 L : 1 M), complex(14); 4.87 g, 0.02 mol of CdSO4·8H2O (1 L : 1 M), complex(15); 6.05 g, 0.02 mol of Hg(OAc)2·4H2O (1 L : 1 M), complex(16); 5.13 g, 0.02 mol of FeCl3·6H2O (1 L : 1 M), complex(17). The reaction mixtures were refluxed with stirring for 2–4 hrs range, depending on the nature of the metal ion and the anion. The precipitates so formed were filtrated off, washed with ethanol, and dried in vacuum over P2O5. Analytical data of the metal complexes are given in (Table 1).


## 2.3. Biological Activity

Evaluation of the cytotoxic activity of the ligand and its metal complexes was carried out in the Pathology Department, Faculty of Medicine, El-Menoufia University, Egypt. The evaluation process was carried out in vitro using the Sulfo-Rhodamine-B-stain (SRB) [25]. Cells were plated in 96-multiwell plate (104 cells/well) for 24 hrs before treatment with the complexes to allow attachment of cell to the wall of the plate. Different concentrations of the compounds under test in DMSO (0, 5, 12.5, 25, and 50 μg/mL) were added to the cell monolayer, triplicate wells being prepared for each individual dose. Monolayer cells were incubated with the complexes for 48 hrs at 37°C and in atmosphere of 5% CO2. After 48 hrs, cells were fixed, washed, and stained with Sulfo-Rhodamine-B-stain. Excess stain was wash with acetic acid and attached stain was recovered with Tris EDTA buffer. Color intensity was measured in an ELISA reader. The relation between surviving fraction and drug concentration is plotted to get the survival curve for each tumor cell line after addition the specified compound.


## 3. Results and Discussion

All complexes are colored, stable at room temperature, nonhygroscopic, partially soluble in common organic solvents such as CHCl3, and appreciably soluble in DMF and DMSO. The analytical and physical data (Table 1) and spectral data (Tables 2–4) are compatible with the proposed structures (Figure 2). Many attempts have been made to grow up single crystal but no diffractable crystals have been grown till now. The molar conductances of the complexes in 10−3 M DMF at 25°C are in the 13.0–33.4 ohm−1 cm2 mol−1 range, indicating a nonelectrolytic nature [26, 27]. The relative high values for some complexes suggest partial dissociation in DMF. The elemental analyses indicated that, all complexes were formed in 1 L : 1 M molar ratio, except complex(3) which is found to be formed in 1 L : 2 M molar ratio.Table 2
IR frequencies of the bands (cm−1) of ligand [H2L], (1) and its metal complexes.
Comp. number
Molecular formula
ν(H2O/OH)
ν(NH)
ν(C=O)acetyl
ν(C=O)amide
ν(C=N)imine
ν(C=N)oxime
ν(N–O)
ν(OAc)/SO4
ν(M–O)
ν(M–N)
ν(M–Cl)
(1)
[H2L]C12H13N3O4
3360, 3315
3215
1700
1664
1613
1587
1147, 1000, 931
—
—
—
—
(2)
[(HL)Cu(OAc)(H2O)2]·H2O   C14H21N3O9Cu
34303520–3080
3260
1680
1650
1605
1575
1155, 1040, 925
1443, 1342
600
575
—
(3)
[(HL)Cu2(OAc)3(H2O)5]·   C18H33N3O15Cu2
34313575–3000
3260
1709
1652
1601
1567
1170, 1066930
1482, 1345
682
582
—
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
34213500–3180
3217
1712
1669
1601
1564
1163, 1068925
1483, 1382
679
580
—
(5)
[(H2L)CuCl2(H2O)]·2H2O   C12H19Cl2N3O7Cu
3434–3380
3141
1695
1659
1606
1533
1154, 11281013, 954
—
662
575
443
(6)
[(H2L)Cu(SO4)(H2O)2]   C12H17N3O10SCu
35003250–3100
3202
1717
1667
1611
1529
1190, 1104977, 915
1260, 1071, 867, 650
618
522
—
(7)
[(HL)Ni(OAc)(H2O)2]·2H2O   C14H23N3O10Ni
34643522–3247
3217
1760
1674
1605
1564
11501024, 909
1456, 1339
682
620
—
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
34283560–3280
3165
1700
1670
1609
1531
1160, 11021043, 925
1240, 1102, 850, 690
592
528
—
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
34953560–3300
3189
1717
1671
1609
1526
1150, 10651022, 920
1465, 1326
624
535
(10)
[(H2L)Co(SO4)(H2O)2]·3H2O   C13H26N3O13SCo
33003560–3215
3215
1700
1652
1610
1553
1152, 1030, 956
1232, 1092858, 650
605
545
—
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
34083570–3300
3217
1695
1600
1595
1565
1170, 11491029, 920
1460, 1338
654
587
(12)
[(HL)Zn(OAc)(H2O)2]·2H2O   C14H23N3O10Zn
33153600–3280
3214
1700
1673
1601
1566
1154, 1039, 985
1523, 1391
597
535
(13)
[(H2L)Zn(SO4)(H2O)2]·3H2O   C12H23N3O13SZn
33803580–3280
3285
1717
1653
1603
1552
1154, 1042, 951
1501, 1366
589
537
(14)
[(HL)Cd(OAc)(H2O)2]·2H2O   C14H23N3O10Cd
33803587–3210
3206
1705
1675
1613
1534
1155, 1044, 979
1237, 1115, 1060, 885, 650
616
526
(15)
[(H2L)Cd(SO4)(H2O)2]·2H2O   C12H21N3O12SCd
34283652–3275
3217
1710
1652
1587
1565
1160, 1014, 956
1500, 1376
612
588
(16)
[(HL)Hg(OAc)]·H2O   C14H17N3O7Hg
33653525–3260
3200
1710
1644
1609
1543
1151, 1008, 936
—
665
524
482
(17)
[(HL)Fe(Cl)2(H2O)]·3H2O   C12H20N3O8Cl2FeTable 3
Electronic spectra (nm) and magnetic moments (BM) for the ligand(1) and metal complexes.
Comp. number
Molecular formula
λ
m
a
x (nm)
μ
e
f
f (BM)
ν
2
/
ν
1
(1)
[H2L]C12H13N3O4
275 nm (ε = 5.86 × 10−4 L mol−1 cm−1)  315 nm (ε = 8.45 × 10−4 L mol−1 cm−1)  350 nm (ε  = 9.20 × 10−4 L mol−1 cm−1)
—
—
(2)
[(HL)Cu(OAc)(H2O)2]·H2O   C14H21N3O9Cu
265, 302, 380, 495, 550, 625
1.79
—
(3)
[(HL)Cu2(OAc)3(H2O)5]·   C18H33N3O15Cu2
268, 308, 380, 425,  550, 626
1.68
—
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
270, 370, 390, 450, 570, 610
1.81
—
(5)
[(H2L)CuCl2(H2O)]·2H2O   C12H19Cl2N3O7Cu
260, 300, 370, 460, 565, 605
1.77
—
(6)
[(H2L)Cu(SO4)(H2O)2]   C12H17N3O10SCu
265, 300, 400, 505, 595, 605
1.78
—
(7)
[(HL)Ni(OAc)(H2O)2]·2H2O   C14H23N3O10Ni
272, 303, 420, 525, 617, 687, 749
2.85
1.06
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
265, 300, 415, 510, 620, 690, 720
2.95
1.04
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
275, 301, 420, 598, 620
5.11
1.03
(10)
[(H2L)Co(SO4)(H2O)2]·3H2O   C13H26N3O13SCo
302, 397, 430, 556, 621, 265
5.52
1.11
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
265, 298, 390, 425, 585, 611
6.10
—
(12)
[(HL)Zn(OAc)(H2O)2]·2H2O   C14H23N3O10Zn
270, 305, 365
Diamagnetic
—
(13)
[(H2L)Zn(SO4)(H2O)2]·3H2O   C12H23N3O13SZn
265, 300, 378
Diamagnetic
—
(14)
[(HL)Cd(OAc)(H2O)2]·2H2O   C14H23N3O10Cd
268, 305, 360
Diamagnetic
—
(15)
[(H2L)Cd(SO4)(H2O)2]·2H2O   C12H21N3O12SCd
265, 302, 370
Diamagnetic
—
(16)
[(HL)Hg(OAc)]·H2O   C14H17N3O7Hg
260, 305, 370
Diamagnetic
—
(17)
[(HL)Fe(Cl)2(H2O)]·3H2O   C12H20N3O8Cl2Fe
265, 302, 330, 475, 541, 615
6.11
—Table 4
ESR data for metal(II) complexes(2)–(5).
Complex
g
∥
g
⊥
g
iso
a
A
∥  (G)
A
⊥ (G)
A
i
s
o
b
(G)
G
c
Δ
E
x
y
(cm−1)
Δ
E
x
z (cm−1)
K
⊥
2
K
∥
2
K
g
∥/A∥ (cm−1)
α
2
β
2
β
1
2
2
β
a
d
2 (%)
(2)
—
—
2.18
—
—
—
—
—
—
—
—
—
—
—
—
—
—
—
(3)
—
—
2.10
—
—
—
—
—
—
—
—
—
—
—
—
—
—
—
(4)
2.31
2.06
2.14
135
15
55
5.2
18181
20202
0.7
0.84
0.86
165
0.76
0.92
1.1
170
72.4
(5)
2.25
2.05
2.12
120
12
48
5
17857
21739
0.63
0.67
0.8
173.1
0.67
0.94
1
182
77.7
a
3
g
i
s
o = (g∥ + 2g⊥).
b
3
A
i
s
o = (A∥ + 2A⊥).
c
G = (g∥-2)/(g⊥ − 2).Figure 2
Proposed structures of the ligand [H2L] and its metal complexes.


### 3.1. Mass Spectra of the Ligand

The mass spectrum of the [H2L], ligand showed the molecular ion peak atm/e 263 amu, confirming its formula weight (F.W. 363). The mass fragmentation patterns observed atm/z = 76, 93, 109, 121, 123, 137, 180, 215, and 263 amu correspond to C6H4, C6H4OH, C6H5O2, C7H5O2, C7H7O2, C7H7NO2, C8H8N2O3, C8H13N3O4, and C12H13N3O4 moieties, respectively, supported the suggested structure of the ligand.


### 3.2.1H-NMR Spectra

The1H-NMR spectrum of the ligand indicated the presence of two peaks at δ = 11.83 and 12.07 ppm assigned to proton of anti- and syn-oxime NOH protons, respectively; these two bands disappeared in the presence of D2O, indicating that these protons are acidic and the hydroxyl group can participate in the coordination with the metal ions. The syn : anti ratio was found to be 1 : 1, indicating that the percentages of free and hydrogen bonded OHs are identical [28, 29]. Signals at δ = 9.73 and δ = 10.90 ppm were assigned to the NH protons [28, 30]. The spectrum showed a set of peaks as multiples in the (7.94–6.81 ppm) range, which were assigned to aromatic protons ring [31]. Peaks which appeared at 2.50 and 2.21 ppm were assigned to acetyl and methyl groups respectively [17, 32]. These signals disappeared upon adding D2O.Zn(II), Cd(II), and Hg(II) complexes(12),(14), and(16) showed similar spectra. The peaks assigned to the oxime protons disappeared, indicating its participation in the metal coordination. A set of multiple peaks corresponding to the aromatic protons were observed in the 6.88–7.50 ppm range. Signals corresponding to acetyl and methyl prortons appeared at 2.50 and 1.91 ppm, respectively [17, 32]; these signals were disappeared upon adding D2O. A new signal was observed around 1.87 ppm, which may be assigned to protons of the coordinated acetate group [33].


### 3.3. IR Spectra

The characteristic infrared spectral data of ligand H2L(1) and its metal complexes are listed in Table 2. The spectrum of the ligand showed characteristics absorption broad bands in 3360–3315 cm−1 range, which are due to intra- and intermolecular hydrogen bonding of OH of the oxime groups with the imino nitrogen and carbonyl oxygen atoms [22, 30]. The medium band at 3215 cm−1 was assigned to (NH) group [29]. The band appearing at 1700 cm−1 was assigned to υ(C=O) band, which is less than the expected value. It is deduced, therefore, that the carbonyl group is involved in hydrogen bondings in the ligand. On the other hand, the υ(C=O) band of the amide group appeared at 1664 cm−1 [29]. The υ(C=N) vibrations (imine and oxime groups) appeared at 1613 and 1587 cm−1, respectively [29, 32]. Two strong bands observed at 1147 and 1000 cm−1 which were assigned to υ(N–O) [22, 34]. The splitting of the υ(N–O) vibration into two bands confirmed the presence of two nonequivalent hydrogen bonding formations whereby the intramolecular type is stronger than the intermolecular type. The bonding mode of the ligand in the metal complexes has been deduced by comparing the IR spectra of the complexes with that of the free ligand. Ir spectra showed that the ligand coordinated through the nitrogen atoms of the imine and the oximato (C=N→O) groups. This mode of bonding is supported by negative shifts in bands of these groups and simultaneous increasing in (N→O) band, appearing in the 1150–1170 cm−1 range [35, 36]. In all complexes except complexes(3) and(4), the strong band assigned to υ(C=O) amid stretching band was shifted to lower frequency, indicating involvement of amide keto oxygen in the metal coordinating. The appearance of two characteristic bands at 1482, 1483, 1500, and 1345, 1382, 1376 cm−1 in spectra of complexes(3),(4), and(16), respectively, were attributed to υasym.(COO-) and υsym.(COO-), respectively, indicating the participation of the acetate oxygen in the complex formation [37]. The mode of coordination of acetate group has often been deduced from the magnitude of the observed separation between the υasym.(COO-) and υsym.(COO-). The separation value (Δ) between υasym.(COO-) and υsym.(COO-) for these complexes were 137, 101, and 124 cm−1 suggesting the coordination of acetate group in a monodentate fashion [38, 39]. In addition, complex(3) showed υ(CO2) at 1560 and 1425 cm−1 due to a bridging acetate group. The chloro complexes(5) and(17) showed new bands at 443 and 482 cm−1, respectively, this band was assigned to ν(M–Cl), whereas sulphate complexes(6),(8),(10),(13), and(15) exhibited new bands in the (1232–1260), (1004–1115), (850–885), and (650–690) cm−1 ranges, these values indicated that the sulphate ion is coordinated to the metal ion in a unidentate chelating fashion [35, 40]. The mode of coordination is supported by presence of additional bands in 620–528 and 682–592 cm−1 regions corresponding to υ(M–N) and υ(M–O) bands, respectively [22, 36, 41].


### 3.4. Electronic Spectra and Magnetic Moments

The electronic absorption spectral data of the ligand and its metal complexes in DMF are listed in Table3. The ligand showed three bands at 275 nm (5.86 × 10−4 L mol−1 cm−1) and 315 nm (ɛ = 8.45 × 10−4 L mol−1 cm−1) and 350 nm (ɛ = 9.2 × 10−4 L mol−1cm−1). The first band may be assigned to π→π∗ transition which is nearly unchanged upon complexation, whereas the second and third bands may be assigned to the n→π∗ and charge transfer transitions of the azomethine and carbonyl groups [30, 42]. These two bands were shifted to lower energy upon complex formation, indicating participation of these groups in coordination with the metal ions. The electronic spectra of copper(II) complexes(2)–(6) were nearly identical and showing bands centered in the 425–505, 550–595, and 605–626 nm ranges assigned to B1g2→A1g2υ1(dx2-y2→dz2),  B1g2→B2g2, υ2(dx2-y2→dxy), and B1g2→Eg2, υ3(dx2-y2→dxy,dyz) transitions, respectively. These transitions indicated that the copper(II) ion has a tetragonally distorted octahedral geometry. This could be due to the Jahn-Teller effect that operates on the d9 electronic ground state of six coordinate system, elongating one trans pair of coordinate bonds and shortening the remaining four ones [43]. The magnetic moments for copper(II) complexes at room temperature were in the 1.68–1.81 range BM, supporting that the complexes have octahedral geometry [44]. The low magnetic moments values of complexes(3) and(4) may be due spin-spin interactions between copper(II) [44]. Nickel(II) complexes(7) and(8) displayed three bands at 420, 617, 749 and 415, 620, 720 nm, respectively, these bands were assigned to A2g3(F)→T2g3(F)υ1, A2g3(F)→T2g3(υ2), and A2g3(F)→T1g3(P)υ3 transitions, indicating octahedral nickel(II) complexes [26, 45]. The υ2/υ1 ratios were 1.06 and 1.04 for complexes(7) and(8), respectively, which are less than the usual range of 1.5–1.75, indicating distorted octahedral nickel(II) complex [26, 45]. The values of magnetic moments for nickel(II) complex(7) and(8) were 2.82 and 2.85 BM, respectively, which are consistent with two unpaired electrons state, confirming octahedral geometry for nickel(II) [45]. The cobalt(II) complexes(9) and(10) exhibited only two bands at (598, 620) and (556, 621) nm ranges, respectively. These bands were assigned to T1g4(F)→T2g4(F)υ1 and T1g4(F)→A2g4(F)υ2 transitions, respectively. The third transition band has not been observed possibly because it is out the range of spectrophotometer (>900 nm). The lower value of υ2/υ1 complex(9) and complex(10) may be due to distortion of the octahedral structures [26]. The magnetic moments for cobalt(II) complexes(9) and(10) at room temperature recorded 5.11 and 5.52 BM, respectively. These values are consistent with high spin cobalt(II) ion (d7). Manganese(II) complex(11) displayed weak bands at 425, 585, and 611 nm. These bands were assigned to A1g6→Eg4, A1g6→T2g4, and A1g6→T1g4 transitions, respectively, corresponding to an octahedral structure for manganese(II) complex [43, 46]. Since all the excited states of Mn(II) ion either quartets or doublets, the absorption spectra of Mn(II) ions have only spin forbidden transitions. Therefore, the intensity of transitions was weak. The value of magnetic moment for manganese(II) complex(11) is 6.10 BM which is consistent with high spin octahedral geometry for manganese(II) [41, 47]. Iron(III) complex(17) showed bands at 475, 541, and 615 nm. The first transition is related to charge transfer from the ligand to iron(III) ion, whereas the other two bands were assigned to A1g6→T1g4 and A1g6→T2g4 transitions, suggesting a distorted octahedral structure around the iron atom [46, 48]. The recorded magnetic moment (6.11 BM) is consistent with the proposed high spin octahedral geometry for iron(III) complex [41, 47]. The observed bands in zinc(II)(12) and(13), cadmium(II)(14) and(15), and mercury(II)(16) complexes (Table 3) are due to interligand transitions within the ligand.


### 3.5. Electron Spin Resonance (ESR)

The ESR spectra of solid copper(II) complexes(2)–(5) at room temperature are characteristic of a species with a d9 configuration and having an axial symmetry type of a d(x2-y2) ground state, which is the most common for copper(II) complexes [22]. Complexes(2) and(3) showed isotropic type with giso=2.18 and 2.10, whereas complexes(4) and(5) showed axial type with g∥>g⊥>2.04, indicating a tetragonal distortion [26, 33], corresponding to elongation along the fold symmetry axis Z. The g-values are related by the expression G=(g∥-2)/(g⊥-2). If G>4.0, then, the local tetragonal axes are aligned parallel or only slightly misaligned; if G<4.0, the significant exchange coupling is present. Complexes(4) and(5) showed G value ≥ 4.0, indicating that tetragonal axes are present. Also, these complexes showed g∥≤2.3, suggesting considerable covalent bond character around the copper(II) ion [49, 50]. Also, the in-plane σ-covalence parameter, α2(Cu), was calculated by(3)α2Cu=A∥0.036+g∥-2.002+37g⊥-2.002+0.04.The calculated values for(4) and(5) are 0.76 and 0.67 (Table 4), suggesting covalent bond character [26, 51]. The g∥/A∥ is taken as an indication for the stereochemistry of the copper(II) complexes. Karlin has suggested that this ratio may be an empirical indication of the stereochemistry of copper(II) complex [52]. The value g∥/A∥ quotient in the (165–173.1) cm−1 range is expected for copper(II) complexes within perfectly square based geometry and those higher than 150 cm−1 for tetragonally distorted complexes. The values for copper(II) complexes(4) and(5) are associated with a tetragonally distorted field around copper(II) centers. For copper(II) complexes with 2B1 ground state, the g-values can be related to the parallel (K∥) and perpendicular (K⊥) components of the orbital reduction factor (K⊥) as follows [51]:(4)K∥2=g∥-2.0023ΔExy8λo,K⊥2=g⊥-2.0023ΔExz2λo,K2=K∥2+2K⊥23,where λo is the spine orbit coupling of free copper ion (−828 cm−1) and ΔExy and ΔExz are the electronic transition energies of 2B1→2B2 and 2B1→2E, respectively. For the purpose of calculation, it was assumed that the maximum in the band corresponds to ΔExy and ΔExz can be taken from the wavelength of these bands. From the above relations, the orbital reduction factors (K∥, K⊥, and K) which are a measure of covalence can be calculated. For an ionic environment, K=1, and for a covalent environment, K<1, the lower the value of K, the greater the covalent character. The values of K for(4) and(5) (Table 4) showed considerable covalent bond character. The in-plane and out-of-plane π-bonding coefficients (β12 and β2), respectively, are dependent upon to values of ΔExy and ΔExz in the following equations [53]:(5)α2β2=g⊥-2.002ΔE2λο,α2β12=g∥-2.002ΔE8λο.Complexes(4) and(5) showed β12 values 1.1 and 1.0, indicating a moderate degree of ionic character in the in-plane π-bonding, while β2 are 0.92 and 0.94, indicating ionic character in the out-of-plane π-bonding [54]. It is possible to calculate the approximate orbital population for d orbital using the following equations [22]:(6)A∥=Aiso-2B1±74Δg∥,ad2=2B2Bo,where 2Bo is the calculated dipolar coupling for unit occupancy of the d orbital. When the data of complexes(4) and(5) are analyzed, the results suggested an orbital population close to 72.4 and 77.7% d-orbital spin density clearly, the orbit of the unpaired electron is a d(x2-y2) ground state [22]. The ESR spectral data, for copper(II) complexes are shown in Table 4. Co(II)(9) and Mn(II)(11) complex showed isotropic spectra with giso=2.1 and 2.003, respectively.


### 3.6. Thermal Analyses (DTA and TGA)

IR spectral data (Table2) indicates the presence of water molecules; thermal analyses were carried out to ascertain their nature, and to give an insight into the thermal stability of the studied compounds. The results showed that there is a good agreement in the weight loss between the calculated and the proposed formulae. The thermal analyses imply that all complexes generally decomposed in several steps (Table 5). The DTA and TGA thermogram of complex(4) showed that the complex decomposed in four steps. The first peak at 120°C with a weight loss of 7.0% (calcd. 7.08%) is assigned to elimination of three hydrated water molecules, which is accompanied by an endothermic peak. The second step appeared as an exothermic peak at 160°C, assigned to loss of two CH3COOH molecules with weight loss 7.81% (calcd. 7.74%). The third step appeared as an exothermic peak at 298°C, referring to melting point of the complex. The fourth step at 475°C with a weight loss of 73.36% (calcd. 74.48) implies completing decomposition of this complex that ended with the formation CuO that is accompanied by an exothermic peak. The TG and DTA thermogram of complex(8) showed that the complex decomposed in four steps. The first peak at 90°C with a weight loss of 3.69% (calcd. 3.60%) assigned to the elimination one hydrated water molecule, which is accompanied by an endothermic peak. The second step at 195°C with a weight loss of 7.48% (calcd. 7.39%) is assigned to elimination of two coordinating water molecules that is accompanied by an endothermic peak. The third step appeared as an exothermic peak at 275°C, assigned to loss of one H2SO4 molecule with weight loss 20.27 (calcd. 20.13%). The fourth step appeared as an exothermic peak at 400°C, referring to melting point of the complex. The fifth step at 460°C with a weight loss of 67.51% (calcd. 68.73) implies the complete decomposition of this complex that ended with the formation NiO that is accompanied by an exothermic peak. The TG and DTA thermogram of complexes(9) and(11) showed that these complexes decomposed in five steps. The first peak appeared at 79 and 90°C with a weight loss of 11.60 and 11.87% (calcd. 11.49 and 11.74%), respectively, is assigned to removal of three hydrated water molecules from each complex; this weight loss is accompanied by an endothermic peak. The second peak at 140 and 170°C with a weight loss of 7.24 and 7.90% (calcd. 7.35 and 7.73%) is assigned to elimination of two coordinated water molecules from each complex; this is accompanied by an endothermic peak. The third step at 170 and 258 associated a weight loss of 12.46 and 12.86% (calcd. 12.25 and 13.04%) is assigned to the elimination of one acetate ion from each complex, this step accompanied with endothermic peak. The fourth step appeared as an exothermic peak at 325 and 319°C, referring to the melting point of the two complexes, respectively. The fifth step at 450 and 445°C with a weight loss of 51.79 and 64.13% (calcd. 52.98 and 65.18) implies to complete decomposition of these complexes, respectively, leaving metal oxide that is accompanied by an exothermic peak:(7)HLCuOAcH2O2·2H2O→60°CHLCuOAcH2O2·2H2OHLCuOAcH2O2→150°CHLCuOAc+2H2OHLCuOAc→265°CLCu+CH3COOHLCu→570°CCuO+organic  residuesTable 5
Thermal analyses for metal complexes.
Comp. number
Molecular formula
Temp. (°C)
DTA (peak)
TGA (Wt. loss %)
Assignments
Endo
Exo
Calc.
Found
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
120
Endo
—
7.08
7.00
Loss of hydrated 3H2O
160
Endo
—
8.33
8.27
Loss of  CH3COOH
240
Endo
—
9.09
9.02
Loss of CH3COOH
298
Endo
—
—
—
Melting point
350
—
Exo
—
—
395
—
Exo
—
—
475
—
Exo
13.47
13.53
Decomposition process with the formation of CuO
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
92
Endo
—
3.69
3.38
Loss of hydrated 1H2O
186
Endo
—
7.67
7.62
Loss of coordinated 2H2O
275
Endo
—
22.17
22.03
Loss of H2SO4
314
Endo
—
—
—
Melting point
400
—
Exo
—
—
420
—
Exo
—
—
460
—
Exo
22.16
22.88
Decomposition process with the formation of NiO
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
79
Endo
—
11.48
11.01
Loss of hydrated 3H2O
140
Endo
—
8.65
8.47
Loss of coordinated 2H2O
170
Endo
—
15.52
15.25
Loss of CH3COOH
325
Endo
—
—
—
Melting point
400
—
Exo
—
—
430
—
Exo
—
—
450
—
Exo
13.39
13.55
Decomposition process with the formation of CoO
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
90
Endo
—
11.58
11.86
Loss of hydrated 3H2O
170
Endo
—
8.73
8.47
Loss of coordinated 2H2O
258
Endo
—
15.65
15.69
Loss of CH3COOH
319
Endo
—
—
—
Melting point
390
—
Exo
—
—
413
—
Exo
—
—
445
—
Exo
12.93
12.71
Decomposition process with the formation of MnO


### 3.7. Cytotoxicity Activity

The cytotoxic activity of the oxime hydrazone ligand H2L(1) and its metal complexes(2),(3),(4), and(9) was evaluated against human liver HepG2 cancer cell, (HepG2 cell line) within 0.1–100 μg/L concentration range. The IC50 values were calculated for each compound and results are presented in Figure 3 and Table 6. As shown, most complexes displayed significantly cytotoxic activities compared to Sorafenib (Nexavar) standard drug. It seems that changing the anion, coordination sites, and the nature of the metal ion has effect on the biological behavior. Cytotoxicity activity of the complexes may be attributed to the central metal atom which was explained by Tweedy’s chelation theory [55, 56]. Cytotoxicity results indicated that all tested complexes (IC50 = 2.24–6.49 μM) (except complex(9) with IC50 = 36.80 μM) demonstrated potent cytotoxicity against HepG2 cancer cells. Copper complex(4) showed the highest cytotoxicity effect with IC50 value of 2.24 μM, followed by complex(3) with IC50 value 2.67 μM and then complex(2) with IC50 value 6.49 μM. It was observed also that all complexes are more active than the free ligand. This indicated enhancing of the antitumor activity upon coordination. The enhancement of cytotoxic activity may be assigned to that the positive charge of the metal increased the acidity of coordinated ligand that bears protons, leading to stronger hydrogen bonds which enhanced the biological activity [57, 58]. It seems that changing the anion, coordination sites, and the nature of the metal ion has a pronounced effect on the biological behavior by altering the binding ability of DNA [59, 60]. Gaetke and Chow had reported that metal has been suggested to facilitate oxidative tissue injury through a free-radical mediated pathway analogous to the Fenton reaction [61].Table 6
Cytotoxic activity (IC50) of the ligand and some metal complexes against human liver   HepG2 cancer.
Number
Compound
IC50 (μM)
Ligand
H2L,  (1)
107 ± 5.2
(2)
[(HL)Cu(OAc)(H2O)2]·H2O
6.49 ± 1.7
(3)
[(HL)Cu2(OAc)3(H2O)5]
2.67 ± 1.2
(4)
[(H2L)2Cu(OAc)2]·3H2O
2.24 ± 1.2
(9)
[(HL)Co(OAc)(H2O)2]·3H2O
36.80 ± 4.0
Standard
Sorafenib (Nexavar)
11.8 ± 3.2Figure 3
IC50 values of the ligand H2L(1) and its metal complexes(2),(3),(4), and(9) against human liver HepG2 cancer cell.


## 3.1. Mass Spectra of the Ligand

The mass spectrum of the [H2L], ligand showed the molecular ion peak atm/e 263 amu, confirming its formula weight (F.W. 363). The mass fragmentation patterns observed atm/z = 76, 93, 109, 121, 123, 137, 180, 215, and 263 amu correspond to C6H4, C6H4OH, C6H5O2, C7H5O2, C7H7O2, C7H7NO2, C8H8N2O3, C8H13N3O4, and C12H13N3O4 moieties, respectively, supported the suggested structure of the ligand.


## 3.2.1H-NMR Spectra

The1H-NMR spectrum of the ligand indicated the presence of two peaks at δ = 11.83 and 12.07 ppm assigned to proton of anti- and syn-oxime NOH protons, respectively; these two bands disappeared in the presence of D2O, indicating that these protons are acidic and the hydroxyl group can participate in the coordination with the metal ions. The syn : anti ratio was found to be 1 : 1, indicating that the percentages of free and hydrogen bonded OHs are identical [28, 29]. Signals at δ = 9.73 and δ = 10.90 ppm were assigned to the NH protons [28, 30]. The spectrum showed a set of peaks as multiples in the (7.94–6.81 ppm) range, which were assigned to aromatic protons ring [31]. Peaks which appeared at 2.50 and 2.21 ppm were assigned to acetyl and methyl groups respectively [17, 32]. These signals disappeared upon adding D2O.Zn(II), Cd(II), and Hg(II) complexes(12),(14), and(16) showed similar spectra. The peaks assigned to the oxime protons disappeared, indicating its participation in the metal coordination. A set of multiple peaks corresponding to the aromatic protons were observed in the 6.88–7.50 ppm range. Signals corresponding to acetyl and methyl prortons appeared at 2.50 and 1.91 ppm, respectively [17, 32]; these signals were disappeared upon adding D2O. A new signal was observed around 1.87 ppm, which may be assigned to protons of the coordinated acetate group [33].


## 3.3. IR Spectra

The characteristic infrared spectral data of ligand H2L(1) and its metal complexes are listed in Table 2. The spectrum of the ligand showed characteristics absorption broad bands in 3360–3315 cm−1 range, which are due to intra- and intermolecular hydrogen bonding of OH of the oxime groups with the imino nitrogen and carbonyl oxygen atoms [22, 30]. The medium band at 3215 cm−1 was assigned to (NH) group [29]. The band appearing at 1700 cm−1 was assigned to υ(C=O) band, which is less than the expected value. It is deduced, therefore, that the carbonyl group is involved in hydrogen bondings in the ligand. On the other hand, the υ(C=O) band of the amide group appeared at 1664 cm−1 [29]. The υ(C=N) vibrations (imine and oxime groups) appeared at 1613 and 1587 cm−1, respectively [29, 32]. Two strong bands observed at 1147 and 1000 cm−1 which were assigned to υ(N–O) [22, 34]. The splitting of the υ(N–O) vibration into two bands confirmed the presence of two nonequivalent hydrogen bonding formations whereby the intramolecular type is stronger than the intermolecular type. The bonding mode of the ligand in the metal complexes has been deduced by comparing the IR spectra of the complexes with that of the free ligand. Ir spectra showed that the ligand coordinated through the nitrogen atoms of the imine and the oximato (C=N→O) groups. This mode of bonding is supported by negative shifts in bands of these groups and simultaneous increasing in (N→O) band, appearing in the 1150–1170 cm−1 range [35, 36]. In all complexes except complexes(3) and(4), the strong band assigned to υ(C=O) amid stretching band was shifted to lower frequency, indicating involvement of amide keto oxygen in the metal coordinating. The appearance of two characteristic bands at 1482, 1483, 1500, and 1345, 1382, 1376 cm−1 in spectra of complexes(3),(4), and(16), respectively, were attributed to υasym.(COO-) and υsym.(COO-), respectively, indicating the participation of the acetate oxygen in the complex formation [37]. The mode of coordination of acetate group has often been deduced from the magnitude of the observed separation between the υasym.(COO-) and υsym.(COO-). The separation value (Δ) between υasym.(COO-) and υsym.(COO-) for these complexes were 137, 101, and 124 cm−1 suggesting the coordination of acetate group in a monodentate fashion [38, 39]. In addition, complex(3) showed υ(CO2) at 1560 and 1425 cm−1 due to a bridging acetate group. The chloro complexes(5) and(17) showed new bands at 443 and 482 cm−1, respectively, this band was assigned to ν(M–Cl), whereas sulphate complexes(6),(8),(10),(13), and(15) exhibited new bands in the (1232–1260), (1004–1115), (850–885), and (650–690) cm−1 ranges, these values indicated that the sulphate ion is coordinated to the metal ion in a unidentate chelating fashion [35, 40]. The mode of coordination is supported by presence of additional bands in 620–528 and 682–592 cm−1 regions corresponding to υ(M–N) and υ(M–O) bands, respectively [22, 36, 41].


## 3.4. Electronic Spectra and Magnetic Moments

The electronic absorption spectral data of the ligand and its metal complexes in DMF are listed in Table3. The ligand showed three bands at 275 nm (5.86 × 10−4 L mol−1 cm−1) and 315 nm (ɛ = 8.45 × 10−4 L mol−1 cm−1) and 350 nm (ɛ = 9.2 × 10−4 L mol−1cm−1). The first band may be assigned to π→π∗ transition which is nearly unchanged upon complexation, whereas the second and third bands may be assigned to the n→π∗ and charge transfer transitions of the azomethine and carbonyl groups [30, 42]. These two bands were shifted to lower energy upon complex formation, indicating participation of these groups in coordination with the metal ions. The electronic spectra of copper(II) complexes(2)–(6) were nearly identical and showing bands centered in the 425–505, 550–595, and 605–626 nm ranges assigned to B1g2→A1g2υ1(dx2-y2→dz2),  B1g2→B2g2, υ2(dx2-y2→dxy), and B1g2→Eg2, υ3(dx2-y2→dxy,dyz) transitions, respectively. These transitions indicated that the copper(II) ion has a tetragonally distorted octahedral geometry. This could be due to the Jahn-Teller effect that operates on the d9 electronic ground state of six coordinate system, elongating one trans pair of coordinate bonds and shortening the remaining four ones [43]. The magnetic moments for copper(II) complexes at room temperature were in the 1.68–1.81 range BM, supporting that the complexes have octahedral geometry [44]. The low magnetic moments values of complexes(3) and(4) may be due spin-spin interactions between copper(II) [44]. Nickel(II) complexes(7) and(8) displayed three bands at 420, 617, 749 and 415, 620, 720 nm, respectively, these bands were assigned to A2g3(F)→T2g3(F)υ1, A2g3(F)→T2g3(υ2), and A2g3(F)→T1g3(P)υ3 transitions, indicating octahedral nickel(II) complexes [26, 45]. The υ2/υ1 ratios were 1.06 and 1.04 for complexes(7) and(8), respectively, which are less than the usual range of 1.5–1.75, indicating distorted octahedral nickel(II) complex [26, 45]. The values of magnetic moments for nickel(II) complex(7) and(8) were 2.82 and 2.85 BM, respectively, which are consistent with two unpaired electrons state, confirming octahedral geometry for nickel(II) [45]. The cobalt(II) complexes(9) and(10) exhibited only two bands at (598, 620) and (556, 621) nm ranges, respectively. These bands were assigned to T1g4(F)→T2g4(F)υ1 and T1g4(F)→A2g4(F)υ2 transitions, respectively. The third transition band has not been observed possibly because it is out the range of spectrophotometer (>900 nm). The lower value of υ2/υ1 complex(9) and complex(10) may be due to distortion of the octahedral structures [26]. The magnetic moments for cobalt(II) complexes(9) and(10) at room temperature recorded 5.11 and 5.52 BM, respectively. These values are consistent with high spin cobalt(II) ion (d7). Manganese(II) complex(11) displayed weak bands at 425, 585, and 611 nm. These bands were assigned to A1g6→Eg4, A1g6→T2g4, and A1g6→T1g4 transitions, respectively, corresponding to an octahedral structure for manganese(II) complex [43, 46]. Since all the excited states of Mn(II) ion either quartets or doublets, the absorption spectra of Mn(II) ions have only spin forbidden transitions. Therefore, the intensity of transitions was weak. The value of magnetic moment for manganese(II) complex(11) is 6.10 BM which is consistent with high spin octahedral geometry for manganese(II) [41, 47]. Iron(III) complex(17) showed bands at 475, 541, and 615 nm. The first transition is related to charge transfer from the ligand to iron(III) ion, whereas the other two bands were assigned to A1g6→T1g4 and A1g6→T2g4 transitions, suggesting a distorted octahedral structure around the iron atom [46, 48]. The recorded magnetic moment (6.11 BM) is consistent with the proposed high spin octahedral geometry for iron(III) complex [41, 47]. The observed bands in zinc(II)(12) and(13), cadmium(II)(14) and(15), and mercury(II)(16) complexes (Table 3) are due to interligand transitions within the ligand.


## 3.5. Electron Spin Resonance (ESR)

The ESR spectra of solid copper(II) complexes(2)–(5) at room temperature are characteristic of a species with a d9 configuration and having an axial symmetry type of a d(x2-y2) ground state, which is the most common for copper(II) complexes [22]. Complexes(2) and(3) showed isotropic type with giso=2.18 and 2.10, whereas complexes(4) and(5) showed axial type with g∥>g⊥>2.04, indicating a tetragonal distortion [26, 33], corresponding to elongation along the fold symmetry axis Z. The g-values are related by the expression G=(g∥-2)/(g⊥-2). If G>4.0, then, the local tetragonal axes are aligned parallel or only slightly misaligned; if G<4.0, the significant exchange coupling is present. Complexes(4) and(5) showed G value ≥ 4.0, indicating that tetragonal axes are present. Also, these complexes showed g∥≤2.3, suggesting considerable covalent bond character around the copper(II) ion [49, 50]. Also, the in-plane σ-covalence parameter, α2(Cu), was calculated by(3)α2Cu=A∥0.036+g∥-2.002+37g⊥-2.002+0.04.The calculated values for(4) and(5) are 0.76 and 0.67 (Table 4), suggesting covalent bond character [26, 51]. The g∥/A∥ is taken as an indication for the stereochemistry of the copper(II) complexes. Karlin has suggested that this ratio may be an empirical indication of the stereochemistry of copper(II) complex [52]. The value g∥/A∥ quotient in the (165–173.1) cm−1 range is expected for copper(II) complexes within perfectly square based geometry and those higher than 150 cm−1 for tetragonally distorted complexes. The values for copper(II) complexes(4) and(5) are associated with a tetragonally distorted field around copper(II) centers. For copper(II) complexes with 2B1 ground state, the g-values can be related to the parallel (K∥) and perpendicular (K⊥) components of the orbital reduction factor (K⊥) as follows [51]:(4)K∥2=g∥-2.0023ΔExy8λo,K⊥2=g⊥-2.0023ΔExz2λo,K2=K∥2+2K⊥23,where λo is the spine orbit coupling of free copper ion (−828 cm−1) and ΔExy and ΔExz are the electronic transition energies of 2B1→2B2 and 2B1→2E, respectively. For the purpose of calculation, it was assumed that the maximum in the band corresponds to ΔExy and ΔExz can be taken from the wavelength of these bands. From the above relations, the orbital reduction factors (K∥, K⊥, and K) which are a measure of covalence can be calculated. For an ionic environment, K=1, and for a covalent environment, K<1, the lower the value of K, the greater the covalent character. The values of K for(4) and(5) (Table 4) showed considerable covalent bond character. The in-plane and out-of-plane π-bonding coefficients (β12 and β2), respectively, are dependent upon to values of ΔExy and ΔExz in the following equations [53]:(5)α2β2=g⊥-2.002ΔE2λο,α2β12=g∥-2.002ΔE8λο.Complexes(4) and(5) showed β12 values 1.1 and 1.0, indicating a moderate degree of ionic character in the in-plane π-bonding, while β2 are 0.92 and 0.94, indicating ionic character in the out-of-plane π-bonding [54]. It is possible to calculate the approximate orbital population for d orbital using the following equations [22]:(6)A∥=Aiso-2B1±74Δg∥,ad2=2B2Bo,where 2Bo is the calculated dipolar coupling for unit occupancy of the d orbital. When the data of complexes(4) and(5) are analyzed, the results suggested an orbital population close to 72.4 and 77.7% d-orbital spin density clearly, the orbit of the unpaired electron is a d(x2-y2) ground state [22]. The ESR spectral data, for copper(II) complexes are shown in Table 4. Co(II)(9) and Mn(II)(11) complex showed isotropic spectra with giso=2.1 and 2.003, respectively.


## 3.6. Thermal Analyses (DTA and TGA)

IR spectral data (Table2) indicates the presence of water molecules; thermal analyses were carried out to ascertain their nature, and to give an insight into the thermal stability of the studied compounds. The results showed that there is a good agreement in the weight loss between the calculated and the proposed formulae. The thermal analyses imply that all complexes generally decomposed in several steps (Table 5). The DTA and TGA thermogram of complex(4) showed that the complex decomposed in four steps. The first peak at 120°C with a weight loss of 7.0% (calcd. 7.08%) is assigned to elimination of three hydrated water molecules, which is accompanied by an endothermic peak. The second step appeared as an exothermic peak at 160°C, assigned to loss of two CH3COOH molecules with weight loss 7.81% (calcd. 7.74%). The third step appeared as an exothermic peak at 298°C, referring to melting point of the complex. The fourth step at 475°C with a weight loss of 73.36% (calcd. 74.48) implies completing decomposition of this complex that ended with the formation CuO that is accompanied by an exothermic peak. The TG and DTA thermogram of complex(8) showed that the complex decomposed in four steps. The first peak at 90°C with a weight loss of 3.69% (calcd. 3.60%) assigned to the elimination one hydrated water molecule, which is accompanied by an endothermic peak. The second step at 195°C with a weight loss of 7.48% (calcd. 7.39%) is assigned to elimination of two coordinating water molecules that is accompanied by an endothermic peak. The third step appeared as an exothermic peak at 275°C, assigned to loss of one H2SO4 molecule with weight loss 20.27 (calcd. 20.13%). The fourth step appeared as an exothermic peak at 400°C, referring to melting point of the complex. The fifth step at 460°C with a weight loss of 67.51% (calcd. 68.73) implies the complete decomposition of this complex that ended with the formation NiO that is accompanied by an exothermic peak. The TG and DTA thermogram of complexes(9) and(11) showed that these complexes decomposed in five steps. The first peak appeared at 79 and 90°C with a weight loss of 11.60 and 11.87% (calcd. 11.49 and 11.74%), respectively, is assigned to removal of three hydrated water molecules from each complex; this weight loss is accompanied by an endothermic peak. The second peak at 140 and 170°C with a weight loss of 7.24 and 7.90% (calcd. 7.35 and 7.73%) is assigned to elimination of two coordinated water molecules from each complex; this is accompanied by an endothermic peak. The third step at 170 and 258 associated a weight loss of 12.46 and 12.86% (calcd. 12.25 and 13.04%) is assigned to the elimination of one acetate ion from each complex, this step accompanied with endothermic peak. The fourth step appeared as an exothermic peak at 325 and 319°C, referring to the melting point of the two complexes, respectively. The fifth step at 450 and 445°C with a weight loss of 51.79 and 64.13% (calcd. 52.98 and 65.18) implies to complete decomposition of these complexes, respectively, leaving metal oxide that is accompanied by an exothermic peak:(7)HLCuOAcH2O2·2H2O→60°CHLCuOAcH2O2·2H2OHLCuOAcH2O2→150°CHLCuOAc+2H2OHLCuOAc→265°CLCu+CH3COOHLCu→570°CCuO+organic  residuesTable 5
Thermal analyses for metal complexes.
Comp. number
Molecular formula
Temp. (°C)
DTA (peak)
TGA (Wt. loss %)
Assignments
Endo
Exo
Calc.
Found
(4)
[(H2L)2Cu(OAc)2]·3H2O   C28H38N6O15Cu
120
Endo
—
7.08
7.00
Loss of hydrated 3H2O
160
Endo
—
8.33
8.27
Loss of  CH3COOH
240
Endo
—
9.09
9.02
Loss of CH3COOH
298
Endo
—
—
—
Melting point
350
—
Exo
—
—
395
—
Exo
—
—
475
—
Exo
13.47
13.53
Decomposition process with the formation of CuO
(8)
[(H2L)Ni(SO4)(H2O)2]·H2O   C13H22N3O11SNi
92
Endo
—
3.69
3.38
Loss of hydrated 1H2O
186
Endo
—
7.67
7.62
Loss of coordinated 2H2O
275
Endo
—
22.17
22.03
Loss of H2SO4
314
Endo
—
—
—
Melting point
400
—
Exo
—
—
420
—
Exo
—
—
460
—
Exo
22.16
22.88
Decomposition process with the formation of NiO
(9)
[(HL)Co(OAc)(H2O)2]·3H2O   C14H25N3O11CO
79
Endo
—
11.48
11.01
Loss of hydrated 3H2O
140
Endo
—
8.65
8.47
Loss of coordinated 2H2O
170
Endo
—
15.52
15.25
Loss of CH3COOH
325
Endo
—
—
—
Melting point
400
—
Exo
—
—
430
—
Exo
—
—
450
—
Exo
13.39
13.55
Decomposition process with the formation of CoO
(11)
[(HL)Mn(OAc)(H2O)2]·3H2O   C14H25N3O11Mn
90
Endo
—
11.58
11.86
Loss of hydrated 3H2O
170
Endo
—
8.73
8.47
Loss of coordinated 2H2O
258
Endo
—
15.65
15.69
Loss of CH3COOH
319
Endo
—
—
—
Melting point
390
—
Exo
—
—
413
—
Exo
—
—
445
—
Exo
12.93
12.71
Decomposition process with the formation of MnO


## 3.7. Cytotoxicity Activity

The cytotoxic activity of the oxime hydrazone ligand H2L(1) and its metal complexes(2),(3),(4), and(9) was evaluated against human liver HepG2 cancer cell, (HepG2 cell line) within 0.1–100 μg/L concentration range. The IC50 values were calculated for each compound and results are presented in Figure 3 and Table 6. As shown, most complexes displayed significantly cytotoxic activities compared to Sorafenib (Nexavar) standard drug. It seems that changing the anion, coordination sites, and the nature of the metal ion has effect on the biological behavior. Cytotoxicity activity of the complexes may be attributed to the central metal atom which was explained by Tweedy’s chelation theory [55, 56]. Cytotoxicity results indicated that all tested complexes (IC50 = 2.24–6.49 μM) (except complex(9) with IC50 = 36.80 μM) demonstrated potent cytotoxicity against HepG2 cancer cells. Copper complex(4) showed the highest cytotoxicity effect with IC50 value of 2.24 μM, followed by complex(3) with IC50 value 2.67 μM and then complex(2) with IC50 value 6.49 μM. It was observed also that all complexes are more active than the free ligand. This indicated enhancing of the antitumor activity upon coordination. The enhancement of cytotoxic activity may be assigned to that the positive charge of the metal increased the acidity of coordinated ligand that bears protons, leading to stronger hydrogen bonds which enhanced the biological activity [57, 58]. It seems that changing the anion, coordination sites, and the nature of the metal ion has a pronounced effect on the biological behavior by altering the binding ability of DNA [59, 60]. Gaetke and Chow had reported that metal has been suggested to facilitate oxidative tissue injury through a free-radical mediated pathway analogous to the Fenton reaction [61].Table 6
Cytotoxic activity (IC50) of the ligand and some metal complexes against human liver   HepG2 cancer.
Number
Compound
IC50 (μM)
Ligand
H2L,  (1)
107 ± 5.2
(2)
[(HL)Cu(OAc)(H2O)2]·H2O
6.49 ± 1.7
(3)
[(HL)Cu2(OAc)3(H2O)5]
2.67 ± 1.2
(4)
[(H2L)2Cu(OAc)2]·3H2O
2.24 ± 1.2
(9)
[(HL)Co(OAc)(H2O)2]·3H2O
36.80 ± 4.0
Standard
Sorafenib (Nexavar)
11.8 ± 3.2Figure 3
IC50 values of the ligand H2L(1) and its metal complexes(2),(3),(4), and(9) against human liver HepG2 cancer cell.


## 4. Conclusions

New copper(II), Nickel, cobalt(II), manganese(II), zinc(II) and cadmium(II), Mercury(II), and iron(II) metal complexes derived from N′-(3-(hydroxyimino)-4-oxopentan-2-ylidene) salicylic hydrazide (H2L) were synthesized. The analytical and physicochemical data confirmed the composition and structure of the newly obtained compounds. The ligand behaved as monobasic tridentate, monobasic bidentate, neutral tetradentate, and neutral tridentate. The complexes adopted distorted octahedral geometry around the metal ion. The ligand and tested complexes showed a high potential cytotoxic activity against growth human liver cancer HepG2 tumor cell lines) compared to Sorafenib (Nexavar) standard drug. All complexes were found to be more active than the free ligand. This indicates enhancing of antitumor activity upon coordination. Cupper complex(4) showed the highest cytotoxic activity with IC50 2.24 μM followed by complex 2.67 μM. These compounds are promising candidates as anticancer agents because of their high cytotoxic activity.



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*Source: 126023-2015-06-25.xml*