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Original article
10 (
2_suppl
); S3816-S3825
doi:
10.1016/j.arabjc.2014.05.019

Synthesis, structural characterization, electrochemical and biological studies on divalent metal chelates of a new ligand derived from pharmaceutical preservative, dehydroacetic acid, with 1,4-diaminobenzene

, , ,
a Chemistry Department, Faculty of Science, Menoufia University, Shebin El-Kom, Egypt
b Pharmaceutical Analysis Department, Faculty of Pharmacy, Damanhour University, Damanhour, Egypt

⁎Corresponding author. Tel.: +20 100 4844589. hmaahmed@yahoo.co.uk (Hytham M. Ahmed)

Received: , Accepted: ,
© The Author(s) 2014
Disclaimer:
This article was originally published by Elsevier and was migrated to Scientific Scholar after the change of Publisher.

Peer review under responsibility of King Saud University.

Abstract

Cobalt(II), nickel(II), copper(II), zinc(II) and cadmium(II) complexes of new 3-acetyl-4-[(4-aminophenyl)amino]-6-methyl-2H-pyran-2-one (HL1) derived from dehydroacetic acid and 1,4-diaminobenzene were prepared and characterized. The structural features were determined from their elemental analyses, 1H, and 13C-NMR spectra, molar conductivities, magnetic moments, IR, UVvis. spectra, thermal analyses (D.T.A. and T.G.A.) and E.S.R. measurements. Their magnetic susceptibility measurements and low conductance data provide evidence for the mono- or dimeric and non-electrolytic nature of the solid complexes. The E.S.R. spectra of copper(II) complexes show axial type symmetry with covalent or ionic bond character. The electrochemical behavior of the complexes in DMF (dimethylformamide) solvent at 298 K was studied. The biological activity of the ligand and its metal(II) complexes was also studied. The obtained complexes showed higher activities than the free ligand in protecting the Egyptian cotton fields from Spodoptera littoralis larvae.

Keywords

Dehydroacetic acid
Complex syntheses
Spectroscopic studies
Conductivity
Thermal analyses
Magnetism
1

1 Introduction

Dehydroacetic acid (DHA) is used mostly as a fungicide and bactericide (Miller et al., 2001). DHA is commonly used in cosmetic products to prevent alteration and degradation of the product formulations. However, this preservative may be harmful to consumers due to its tendency to induce allergic contact dermatitis. DHA has been widely used as antimicrobial agents in pharmaceuticals and cosmetics because of its broad antimicrobial spectrum with good stability and non-volatility (Kabara, 1984). Hence, a selective reaction for DHA was studied in order to establish a qualitative and quantitative method of analysis for it through the formation of different metal chelates of DHA new ligand.

In recent years, the design, synthesis and structural characterization of Schiff base complexes derived from dehydroacetic acid continue to be a subject of current interest, not only because they can be useful models for understanding the nature of complexes of biological relevance (Que, 1988) but also due to their interesting structural and magnetic properties (Kahn, 1993; Tuna et al., 2000). Transition metal complexes of dehydroacetic acid have shown antifungal properties more active than either dehydroacetic alone or the pure inorganic salts (Miyakado et al., 1982; Rao et al., 1978). Copper(II) and nickel(II) complexes of unsymmetrical Schiff base containing dehydroacetic acid have been prepared and characterized (Cindric et al., 2004). Cobalt(II, III), nickel(II) and copper complexes of dehydroacetic acid N4-dialkyl and 3-azacyclothiosemicarbazones have been studied (Chalaca et al., 2002). Manganese(III) complexes of tetra dentate Schiff bases containing dehydroacetic fragments have been prepared and thoroughly characterized (Fernandez et al., 2002). In this paper, synthesis and structural characterization of Co(II), Ni(II), Cu(II), Zn(II) and Cd(II) of 3-[(4-aminophenyl)amino]-6-methyl-2,4-dioxo-3,4-dihydro-2H-pyran were done.

2

2 Experimental

All chemicals and solvents were reagent grade commercial materials and used as received. C, H, N and Cl were determined at the Analytical Unit of Cairo University, Egypt. Metal contents were determined by standard methods. All complexes were dried in vacuum over P2O5. IR spectra as KBr pellets were recorded on a Perkin-Elmer 681 spectrophotometer. Electronic spectra (solid state) were recorded on a Perkin-Elmer 550 spectrometer. 1H- and 13C-NMR spectra were obtained on a Bruker AM300 MHz instrument with Me4Si as internal reference. Magnetic susceptibilities were measured at 25 °C by the Gouy method with mercuric tetrathiocyanatocobaltate(II) as the magnetic susceptibility standard. Diamagnetic corrections were made using Pascal’s constant. Magnetic susceptibilities were measured at 25 °C from the equation: μeff = 2.88 (χn × T)1/2. Molar conductance values were measured on a tacussel-type CD6NG conductivity bridge using 10−3 M DMF solutions. Thermal analysis (D.T.A. and T.G.A.) was carried out in air (20–800 °C) using a Shimadzu DT-30 thermal analyzer. Cyclic voltammetry was measured in 2 × 10−3 M DMF solution containing 0.1 M of LiCl as supporting electrolyte, using a gold working electrode. E.S.R. measurements (solid state) at room temperature were carried out using a Varian E-109, X-band spectrometer. The spectrometer was calibrated using a standard sample (DPPH). TLC of all compounds confirmed their purity.

2.1

2.1 Preparation of the ligand (HL1)

The ligand (HL1) (1) was prepared by refluxing an equimolar amount of dehydroacetic acid (2 g, 0.008 mol) and 1,4-diaminobenzene (1.73 g, 0.008 mol) in 100 cm3 of ethanol for 1 h, the yellow product obtained was filtered off, washed several times with ethanol and then dried over P4O10.

2.2

2.2 Preparation of metal(II) complexes

Metal(II) complexes were prepared by mixing a stoichiometric ratio (1:1) of the ligand (2.0 g) dissolved in ethanol (50 cm3) to appropriate salt dissolved in ethanol (30 cm3) in the absence of sodium acetate, (2), (7),(8), (9) and (10), and in the presence of sodium acetate (3), (4), (5), (6), (11), (12) and (13). The mixture was heated at reflux for 1–4 h. On cooling to room temperature and reducing the volume, a precipitate appeared which was filtered off, washed with ethanol and subsequently dried over P4O10.

2.3

2.3 Biological application

Spodoptera littoralis larvae were collected from cotton fields in Menoufia Province and reared on castor bean leaves, Ricinus communis, that renewed daily and maintained under laboratory conditions of temperature (27 ± 2 °C) and humidity (65 ± 5% R.H.). Experiments were carried out on the 4th-instar larvae. The chemicals were weighted and dissolved in DMF as the solvent to get the tested concentrations (25, 50, 75 and 100 ppm). The castor bean leaves were tipped for 10 s in the tested concentrations and then introduced to the larvae for 48 h. New untreated castor bean leaves were replaced by treated ones till pupation flour replicates contained 25 larvae one jar were used for each treatment and also the check experiments which carried out by the solvent only. These tests were carried out to define the effect of the tested chemical on the larval mortality.

3

3 Results and discussion

3.1

3.1 General data

All the synthesized compounds are colored solid, stable at room temperature, non-hygroscopic, insoluble in water and partially soluble in common organic solvents such as CHCl3, but soluble in DMF and DMSO. The analytical and physical data (Table 1) and spectral data (Tables 2 and 3) are in agreement with the suggested structures (Fig. 1). The molar conductivities of solutions of metal complexes (10−3 M in DMF) reveal that, all the complexes are non-electrolytes (Geary, 1971). Metal complexes are prepared by direct reaction between the ligand and the appropriate salts with stoichiometric ratio (1:1) in ethanol. Scheme 1 illustrates that, the composition of the metal complexes depends on the nature of the metal salts, the time and the medium of the reaction as shown below.

Table 1 Analytical and physical data of the ligand (1) and its metal(II) complexes.
No. Compound Formulation Color MP (°C) Yield (%) Λa μeff (B.M.) Found (calcd.) (%)
C H N Cl M
(1) HL1 C14H14N2O3 Pale yellow 180 82 – – 64.77 5.57 10.45 – –
(65.1) (5.46) (10.85)
(2) [Co(HL1)(L1) (OAc)(H2O)] C30H32N4O9Co Brown 198 72 5.0 4.2 (56.06) 4.60 8.60 – 9.10
55.3 (4.90) (8.59) (9.0)
(3) [Co(L1)2(H2O)2] C28H30N4O8Co Brown 176 75 0.4 4.3 54.7 4.90 10.93 – 9.20
(54.99) (4.94) (9.16) (9.60)
(4) [Ni(HL1)(OAc)2(H2O)2]0.5C2H5OH C19H27N2O9.5Ni Dark green 200 73 3.2 3.1 46.00 4.65 5.51 – 12.8
(46.2) (5.5) (5.66) (11.87)
(5) [Ni(HL1)(Cl)2(H2O)2]·H2O C14H20N2O6Cl2Ni Pale Brown 300 70 3.4 3.2 37.40 4.70 5.4 15.2 12.6
(38.0) (4.56) (6.33) (16.0) (13.1)
(6) [Ni(HL1)(NO3)2(H2O)2]·3H2O C14H24N4O14Ni Brownish orange >300 76 1.2 3.3 31.10 4.50 10.63 – 10.8
(31.66) (4.56) (10.57) (11.1)
(7) [Cu(HL1)(OAc)2]·0.5H2O C18H21N2O7.5Cu Dark green 178 75 0.9 2.0 47.60 4.90 6.10 – 13.7
(48.16) (4.68) (6.24) (14.16)
(8) [Cu(L1)(Cl)(H2O)]2 C28H30N4O8Cl2Cu2 Dark Brown 170 68 4.6 1.41 44.80 3.12 6.10 9.10 16.5
(44.92) (4.01) (7.49) (9.49) (16.98)
(9) [Cu(HL1)(NO3)2] C14H14N4O9Cu Green >300 70 1.8 2.1 37.50 3.11 12.00 – 13.8
(37.71) (3.16) (12.57) (14.2)
(10) [Cu(HL1)(Br)2(H2O)]2·2H2O C28H36N4O10Br4Cu2 Black 200 65 2.4 1.4 32.50 3.00 5.55 – 11.8
(32.46) (3.47) (5.41) (12.3)
(11) [Zn (HL1)2(Cl)2]·2.5H2O C28H33N4O8.5 Cl2Zn Brownish red 204 68 4.6 Diamag 47.0 4.10 6.06 – –
(48.18) (4.7) (8.03)
(12) [Cd(HL1)2(Cl)2]·3H2O C28H34N4O9 Cl2Cd Yellow 200 70 0.8 Diamag 44.3 4.70 7.10 9.40 –
(44.6) (4.51) (7.43) (9.00)
(13) [Cd(HL1)(L1)(NO3)(H2O)] C28H29N5O10Cd Buff 190 65 1.2 Diamag 47.7 4.20 10.1 – –
(47.5) (4.09) (9.89) – –
a ohm−1 cm2 mol−1.
Table 2 IR spectral (cm−1) assignments of ligand (1) and its metal(II) complexes.
Compound Assignment (cm−1)
ν(OH)/H2O ν(NH2) ν(NH) ν(C⚌O) ν(C⚌C) ν(M–O) ν(OAc) ν(NO3) ν(M–X)a ν(M–N)
(1) 3667–3219 3461, 3354 3219 1698 1565 – – – – –
1631
(2) 3600–3090 3431, 3356 3231 1695 1579 625 1505 – – 526
1666 1412
(3) 3600–3068 3459, 3354 – 1695 1571 655 – – – 539
1640
(4) 3600–3100 3369,3260 3130 1694 1583 672 1500 – – 527
1660 1415
(5) 3580–3110 3381,3315 3241 1693 1589 633 – – 335 522
1666
(6) 3610–3110 3376, 3320 3257 1689 1586 609 – 1414, 1317 – 536
1652 841, 700
(7) 3631–3075 3427, 3331 3304 1696 1588 680 1548 – – 528
1655 1433
(8) 3610–3150 3444, 3310 – 1695 1564 640 – – 325 520
1632
(9) – 3450, 3284 3242 1673 1554 624 – 1417,1383 – 536
1644 850, 766
(10) 3575–3136 3421,3337 3251 1661 1522 644 – – 300 535
1622
(11) 3590–3129 3461, 3351 3214 1695 1567 634 – – 311 534
1663
(12) 3632–3090 3462, 3354 3217 1697 1568 638 – – 333 533
1631
(13) 3613–3090 3462, 3356 3221 1698 1567 637 – 1465, 1381 – 530
1629 834, 710
a Cl or Br.
Table 3 Electronic spectra of ligand (1) and its metal complexes.
Compound Medium λmax (nm)
(1) DMF 320 (ε = 5.16 × 103 mol−1 cm−1)
370 (ε = 6.0 × 103 mol−1 cm−1)
(2) N.M.a 780, 500, 380, 340, 300, 280
(3) N.M. 730, 525, 380, 340, 310, 280
(4) N.M. 765, 510, 370, 350, 310
(5) N.M. 735, 520, 360, 325, 310
(6) N.M. 710, 520, 355, 325
(7) N.M. 720, 530, 420, 380, 300, 250
(8) N.M. 808, 650, 405, 324, 270
(9) N.M. 720, 500, 420, 380, 350, 280
(10) N.M. 670, 470, 410, 330, 280
(11) N.M. 410, 380, 320, 280
(12) N.M. 410, 350, 310, 250
(13) N.M. 490, 410, 340, 310, 250
a N.M. = Nujol mull.
The proposed structures of HL1 ligand and its metal(II) complexes.
Figure 1
The proposed structures of HL1 ligand and its metal(II) complexes.
Synthetic protocol for metal (II) complexes of HL1 ligand.
Scheme 1
Synthetic protocol for metal (II) complexes of HL1 ligand.

3.2

3.2 1H and 13C-NMR spectrum

The 1H-NMR spectrum of the ligand (1) showed two peaks at 5.3 and 5.8 ppm, attributed to NH and NH2 protons respectively(Xinde et al., 1991; Krishna et al., 1990).The aromatic protons appear as multiple signal in the 7.3–6.5 ppm range. However, the methyl protons appear at 2.1 and 2.4 ppm (Raman et al., 2004).The 13C-NMR spectrum data of the ligand in DMSO shows peaks at 184.5 and 177.3 ppm, assigned to C2⚌O and C4⚌O, respectively, however, the C⚌C, C—N and CH3 peaks appear at 163.5, 144.4 and 40.1 ppm, respectively. The peaks observed in the 129.7–97.4 ppm range are due to the aromatic carbon atoms (Kasumov and Köksal, 2003; Raman et al., 2004).

3.3

3.3 IR spectra

The infrared spectra of the ligand (1) and its metal(II) complexes (Table 2) could be taken as a diagnostic of the mode of the coordination of the ligand to the metal ions. Assignments were made based on typical group frequencies (Xinde et al., 1991). The spectrum of the ligand showed two broad bands in the 3667–3219 and 2924–2028 cm−1 ranges, indicating two types of hydrogen bonded NH groups which could be intra- and inter- molecularly hydrogen bonded to the C⚌O group, Thus, the higher frequency band at 3667–3219 cm−1, is associated with a weaker hydrogen bonding compared with the lower frequency band at 2924–2028 cm−1 of the relatively stronger hydrogen bonding. The spectra of the ligand and its metal(II) complexes showed strong bands in the 3461–3260 and 3304–3130 cm−1 ranges assigned to NH2 and NH groups, respectively (Xinde et al., 1991; Bindu et al., 1998). The ν(C2⚌O)ring, ν(C4⚌O)ring and ν(C⚌C)Ar vibrations of the ligand appeared at 1698, 1631 and 1565 cm−1, respectively (El-Tabl, 2002). The appearance of a strong and broad absorption band (Table 2) in the 3460–3068 cm−1 range of all complexes except 9 indicated the presence of water molecules (Xinde et al., 1991; El-Tabl and Ayad, 2003; Berger et al., 1999). Also, the infrared spectra of the complexes showed that, the ν(NH2) were not affected by complexation, indicating that the terminal NH2 group was not involved in the coordination to the metal ion (Emam, 1988). The shift of ν(NH), (3304–3130) cm−1 in all complexes except 3 and 8, and ν(C4⚌O), (1666–1622) cm−1 bands were compared to the free ligand (Table 2) and indicated coordination of these groups to the metal ions (Raman et al., 2004). The ν(NH) bands disappeared in complexes 3 and 8 and the shift of ν(C4⚌O) bands were compared to the free ligand and indicated coordination of deprotonated of imino group and (C4⚌O) to the metal ion. A band appears at 1589–1522 cm−1 range, assigned to the vibration of the ν(C⚌C)Ar (Athappan and Rajagopal, 1997). The infrared spectra of the acetato complexes 2, 4 and 7 showed new bands in the 1548–1500 and 1433–1412 cm−1 ranges, which were attributed to νa(COO) and νs(COO) respectively (El-Tabl, 2002; Nakamoto, 1978). However, complexes 6, 9 and 13 showed bands at 1465–1414, 1383–1317, 850–834 and 766–700 cm−1 ranges, which were characteristic of a monodentate nitrate group (El-Tabl, 2002; Ma and Angelici, 1980). The presence of a moderate intensity band in the 335–300 cm−1 range for the complexes 5, 8, 10, 11 and 12 corresponded to ν(M—Cl) or ν(M—Br) vibrations (El-Tabl, 2002; Raman et al., 2004; Nakamoto, 1986). Also, the spectra of the metal complexes showed some new bands in the 680–624 and 539–520 cm−1 ranges, which were probably due to the ν(M—O) and ν(M—N) vibrations, respectively (El-Tabl, 2002; Raman et al., 2004).

3.4

3.4 Electronic spectra

The electronic spectra of the ligand in DMF solution and its metal(II) complexes in Nujol mull are summarized in Table 3. The spectrum of the ligand in DMF showed two bands at 320 nm (ε = 5.16 × 103 mol−1 cm−1) and 370 nm (ε = 6.0 × 103 mol−1 cm−1) which might be assigned to the π → π∗ and n → π∗ transitions, respectively (El-Tabl, 2002; Bose et al., 1973). In Nujol mull, the nickel(II) complexes 4, 5 and 6 exhibited two well- defined bands in the 710–765 and 510–520 ranges, which were assigned to the 3T1g(F) ← 3A2g and 3T1g(P) ← 3A2g transitions respectively in an octahedral structure (El-Tabl, 2002; Sallam, 2005). Copper(II) complexes 7 and 9 show a broad absorption band ∼720 nm that strongly favors the square-planar geometry, this peak is attributed to 2A1g ← 2B1g transition(Sallam, 2005; Gao et al., 1997). The broadness of the band could be taken as an indication of distortion from perfect planar geometry. Complex (8) exhibited two absorption bands at 808 and 650 nm, which were attributed to 2T2g(D) ← 2Eg transition which was characterized by square pyramidal geometry of Cu(II) complexes (Lever, 1968). However complex (10) exhibits two bands at 670 and 470 nm, assigned to 2A1g ← 2B1g and 2Eg ← 2B1g transitions which indicated a tetragonal distorted octahedral geometry(Sallam, 2005; Gao et al., 1997). Cobalt(II) complexes 2 and 3 showed two d–d bands at 780, 730 and 500, 525 nm, which could be assigned to 4A2g(P) ← 4T1g(F) (ν2) and 4T2g(P) ← 4T1g(F) (ν3), transitions, respectively, which indicated octahedral geometry for these complexes (Sallam, 2005; Gao et al., 1997). However, complexes 11, 12 and 13 show bands at 410–420 and 340–380 nm ranges, due to intraligand transitions (Table 3).

3.5

3.5 Magnetic moments

The magnetic moments for the metal complexes are shown in Table 1. The values for nickel(II) complexes 4, 5 and (6), were fallen within the 3.1–3.3 BM range, which indicated octahedral geometry around the nickel(II) ion (El-Tabl, 2002; Sallam, 2005). The values of magnetic moment of copper(II) complexes 7 and 9 were 2.0 and 2.1 BM, respectively. This could be taken as an evidence for the presence of square planar copper(II) complexes (Yilmaz et al., 2004). These values indicated that no metal–metal interactions took place between the copper(II) ions, however, complexes 8 and 10 showed values 1.41 and 1.40 BM, respectively, which were less than the spin only magnetic moment (1.73 BM), indicated square pyramidal and tetragonal distorted octahedral arrangement around copper(II) ion respectively. The low magnetic moment values indicated that metal–metal interactions took place between the copper(II) ions, which were in agreement with the previous published data (Raman et al., 2004). The cobalt(II) complexes 2 and 3 showed magnetic values 4.2 and 4.3 BM, which were in good agreement with those observed for most high-spin octahedral cobalt(II) complexes (Sallam, 2005; Figgis and Lewis, 1964; El-Tabl, 1996). However, complexes (11), (12) and (13) showed diamagnetic property.

3.6

3.6 E.S.R. spectra

The E.S.R. spectra of solid copper(II) complexes 7, 8, 9 and 10 at room temperature were characteristic for a d9, system with an axial type symmetry of d ( x 2 - y 2 ) ground state, which was the most common for copper(II) complexes (El-Tabl, 1997; El-Tabl, 1998; El-Tabl and Kashar, 1998). Complexes 7 and 9 showed spectra with g|| > g⊥ > 2.003, which indicated square planar geometry around the copper(II) ion (Sallam, 2005; Procter et al., 1968). However, complexes 8 and 10 showed spectra with g|| > g⊥ > 2.05, which indicated five or six coordinated arrangement around the copper(II) ion in a square pyramidal or distorted octahedral geometry (Sallam, 2005). The E.S.R. parameters are shown in Table 5. The g-values were related by the expression (El-Tabl and Abou-Sekkina, 1999; Kivelson and Neiman, 1961), G = (g|| − 2)/(g⊥ − 2). If G > 4.0, then local tetragonal axes were aligned parallel or slightly misaligned. When G < 4.0, significant exchange coupling was present. Complexes 7 and 9 showed G > 4.0, which indicated the presence of a tetragonal distorted copper(II) complexes, however, complexes 8 and 10 showed G < 4.0, which indicated significant exchange coupling which confirmed by the magnetic moment values (Table 1). Kivelson and Neiman (1961) noted that, for an ionic environment g|| > 2.3 and for a covalent environment g|| < 2.3. The reported theoretical work (Smith, 1970) seemed to confirm this view. The g-values for the complexes 8 and 9 showed covalent bond character, however, complexes 7 and 10 showed ionic bond character (Chan et al., 1995; El-Tabl et al., 2007; Ray and Kauffman, 1990). The g-values of the copper(II) complexes with 2B1g ground state (g|| > g⊥) might be expressed by Ray and Kauffman, 1990

(1)
g ∥ = 2.002 - ( 8 K ∥ 2 λ ° / Δ E xy )
(2)
g ⊥ = 2.002 - ( 2 K ⊥ 2 λ ° / Δ E xz ) where K|| and K⊥ are the parallel and perpendicular components, respectively of the orbital reduction factor (K), λ° is the spin–orbit coupling constant for the free copper and ΔExy and ΔExz are the electron transition energies of 2B2g ← 2B1g and 2Eg ← 2B1g. From the above relations, the orbital reduction factors (K||, K⊥, K) which are a measure of the covalency (Ayad and El-Tabl, 1999) could be calculated. For an ionic environment, K ⩾ 1, and for a covalent environment K < 1, the lower the value of K, the greater is the covalent character. Rearrangement of (1) and (2)
(3)
K ⊥ 2 = ( g ⊥ - 2.002 ) Δ E xz / 2 λ o
(4)
K ∥ 2 = ( g ∥ - 2.002 ) Δ E xy / 8 λ o
(5)
K 2 = ( K ∥ 2 + 2 K ⊥ 2 ) / 3 The value of K listed in Table 5 for the copper(II) complexes indicated that the bonds had covalent character for complexes 8 and 9 and ionic character for complexes 7 and 10 (Ayad and El-Tabl, 1999; Emam, 1988; Masood et al., 1994).
Table 4 Thermal analyses for some metal complexes of ligand (1).
Comp. No. D.T.A. Assignment T.G.A. found (calcd) (%) Loss type
Exo Endo
− 24–77 Removal of alcohol 5.31 (4.63) 0.5C2H5OH
− 182–245 Elimination of coordinated water and partial decomposition 16.26 (16.07) 2H2O + CO2
(4) − 282–336 Partial of degradation 15.43 (14.86) Acetate and ligand
− 379–496 Final degradation and decomposition 47.9 Ligand
− 24–139 Dehydration process 5.88 (5.99) 1H2O
− 139–181 Elimination of coordinated H2O 3.91 (3.99) 1H2O
(5) − 181–367 Elimination of coordinated water and removal of halogens 20.8 (21.0) H2O + 2Cl
− 367–600 Final decomposition 42.04 Ligand
− 24–93 Dehydration process 5.07 (5.08) 1.5H2O
− 143–200 Dehydration process 5.07 (5.08) 1.5H2O
(6) − 200–288 Elimination of coordinated water 6.72 (6.78) 2H2O
288–409 − Partial decomposition 23.51 (23.55) 2NO3
409–504 − Final decomposition 36.12 Ligand
− 165–186 Removal of coordinated H2O 2.57 (2.52) 1H2O
(13) − 190–317 Removal of nitrate and partial dissociation 46.11 NO3
371–581 − Final degradation 32.65 Ligand
Table 5 The E.S.R. parameters of the solid copper(II) complexes of ligand 1.
Compounds g|| g⊥ gisoa Gb ΔExy (cm−1) ΔExz (cm−1) K⊥ K|| K
(7) 2.35 2.07 2.16 5.0 17857 21053 1.00 0.99 1.00
(8) 2.12 2.06 2.08 2.0 15873 20000 0.53 0.84 0.73
(9) 2.27 2.05 2.12 5.4 17857 20202 0.85 0.77 0.73
(10) 2.34 2.1 2.18 3.4 15385 20202 0.89 1.09 1.06
a giso = (g|| + 2g⊥)/3.
b G = (g|| − 2)/(g⊥ − 2).

3.7

3.7 Thermal studies

The D.T.A. and T.G.A. data (Fig. 2) of some metal complexes are given in Table 4. The complexes 4, 5 and 6 decomposed in four steps. The first stage took place in the 24–77, 24–139 and 24–200 °C ranges, corresponding to the removal of 0.5 mol of ethanol (desolvation), release of one and three water molecules (dehydration). The mass losses obtained in this step were 5.31(4.63%) and 5.88(5.99%) for complexes 4 and 5, respectively. The desolvation and dehydration process were associated with endothermic peaks. Dehydration of complex (6) took place in two steps as shown in the D.T.A curve which showed two endothermic peaks in the temperature ranges, 24–93 and 143–200 °C, corresponding to the mass loss of 5.07(5.08%), respectively, due to the removal of three coordinated water. This showed that the presence of a loosely bound lattice water molecule and more strongly bound one (Rao and Reddy, 1999). The second decomposition stage of the complexes 4, 5 and 6 took place at 182–245, 139–181 and200–288 °C ranges which corresponded to the elimination of coordinated water molecule. The mass loss in complex 4, 16.26(16.07%) was attributed to the removal of two molecules of coordinated water and evaporation of CO2 molecule. This step showed weight losses of 3.91(3.99) and 6.72(6.78%) due the removal of one and two molecules of coordinated water in complexes 5 and 6, respectively. The third and the fourth decomposition of complex 4 took place in two steps at 282–336 and 379–496 °C ranges corresponded to elimination of acetate, partial degradation of ligand and final decomposition, respectively. These steps corresponded to mass losses of 15.43(16.07%) and 47.9% associated with two endothermic peaks. The third and fourth decomposition steps of complex 5 took place in 181–367 and 367–600 °C ranges respectively. The observed mass lossses were 20.8(21.0%) and 42.04% due to removal of another molecule of coordinated water associated with elimination of halogens and final decomposition, respectively, as shown in D.T.A. Curves. The third decomposition stage of complex 6 took place in 288–409 °C range. The observed mass loss was 23.51(23.55%) which could be attributed to the removal of volatile nitrates. This process was accompanied by exothermic peak and followed also by a sharp exothermic peak in 409–504 °C range which corresponded to the mass loss of 36.12% attributed to final decomposition. The D.T.A. curve of complex 13 showed an endothermic peak within the temperature range 165–186 °C. This peak was assigned for the elimination of one molecule of coordinated water and it was confirmed by the weight loss of 2.57(2.52%) which observed by the T.G.A. curve. The D.T.A. curve of it also showed another endothermic peak of melting followed by partial dissociation due to removal of coordinated nitrate within the temperature range 190–371 °C which was followed by a steady mass loss and may be due to final degradation together with volatilization of the residue.

A, B, C and D are T.G.A. and D.T.A. diagrams of complexes 4, 5, 6 and 13, respectively (----- T.G.A. and ––––D.T.A.).
Figure 2
A, B, C and D are T.G.A. and D.T.A. diagrams of complexes 4, 5, 6 and 13, respectively (----- T.G.A. and ––––D.T.A.).

3.8

3.8 Electrochemical studies

Cyclic voltamograms of the ligand (1) and its metal complexes were recorded in DMF/MeCN of 50/50 V solution using 0.1 M LiCl as the supporting electrolyte in potential range between −1.08 and 1.00 V. The scan rate was 100 mV s−1. The ligand HL (1) was electrochemically inactive in potential range between −1.8 and 1.0 V. The electrochemical data are presented in Table 6. The nickel(II) complexes 4, 5 and 6 exhibited two redox couples in the negative potential range. The anodic peak (Epa) was in the (−0.97) to (−0.83) V range and the associated cathodic peak (Epc) was in the (−0.88) to (0.60) V range. These data corresponded to two perfectly reversible one electron processes with E1/2 values in the (−0.92) to (−0.71) V range, and with ΔEp in (0.27) to (0.08) V range, Table 6. The first redox wave observed at the less negative potentials corresponded to the Ni ( II ) → - e Ni ( III ) redox couple and the second redox wave observed at more negative potentials corresponds to the Ni ( III ) → + e Ni ( II ) redox couple(Raman et al., 2004; Bindu and Kurup, 1997). The cobalt(II) complexes 2 and 3 were oxidized from Co ( II ) → - e Co ( III ) at Epa values −0.89 and −1.05 V. The reduction peak for Co ( III ) → + e Co ( II ) appeared at −0.80 and −1.0 V with E1/2 values in −0.84 and −1.02 V and ΔEp in 0.09 and 0.05 V ranges, as shown in Table 6. The absence of counter reduction wave for Co ( II ) → + e Co ( I ) may be understood in terms of increasing difficulty to reduce Co(II) in the form of complex and in non-aqueous conditions(Krishnapriya and Kandaswamy, 2005). The copper(II) complexes 8 and 9 showed a well defined reversible redox peak corresponded to the formation of the peak Cu ( II ) → - e Cu ( III ) at Epa values −1.12 and −0.95 V and the associated cathodic peak for Cu ( III ) → + e Cu ( II ) at Epc values −0.85 and −0.77 V with ΔEp(V) 0.27 and 0.18 V and E1/2 values were −0.98 and −0.86 V, respectively [12,45]. However, complex 7 exhibited one irreversible peak characteristic for Cu ( II ) → - e Cu ( III ) at Epa value −0.94 V (Raman et al., 2004; Bond et al., 1988; Zhu et al., 1997). The complexes 11, 12 and 13 showed Epa peaks at −0.76, −0.94 and −0.87 V due to Zn ( II ) → - e Zn ( III ) and Cd ( II ) → - e Cd ( III ) , respectively, however, The Epc peaks appearing at −0.55, −0.65 and −0.67 V, corresponded to Zn ( III ) → + e Zn ( II ) and Cd ( III ) → + e Cd ( II ) , respectively with ΔEp values which were 0.21, 0.29 and 0.20 V and E1/2 −0.65, −0.79 and −0.77 V indicating a reversible process with one electron transfer (Raman et al., 2004; Zhu et al., 1997; Murukan and Mohanan, 2006).

Table 6 Electrochemical data for the metal complexes of ligand 1 in DMF/MeCN solution containing 0.1 M LiCl as a supporting electrolyte at scan rate 100 mV s−1 in potential range −1.8 V to 1.0 V, at room temperature.
Complexes Epa (V) Epc (V) ΔEp (V) E1/2
(2) −0.89 −0.80 0.09 0.84
(3) −1.05 −1.0 0.05 1.02
(4) −0.96 −0.88 0.08 0.92
(5) −0.97 −0.70 0.27 0.83
(6) −0.83 −0.60 0.23 0.71
(7) −0.94 − − −
(8) −1.12 −0.85 0.27 0.98
(9) −0.95 −0.77 0.18 0.86
(11) −0.76 −0.55 0.21 0.65
(12) −0.94 −0.65 0.29 0.79
(13) −0.87 −0.67 0.2 0.77

3.9

3.9 Biological activity of some complexes on the cotton leaf worm Spodoptera littoralis

The Egyptian cotton leaf worm S. littoralis is considered as one of the most destructive cotton pests in Egypt and many other parts in the world. It is a polyphagous, occurs throughout the year, feeds on a great variety of crops and attacks numerous types of plants (Abdel-Hafew, 1978). It causes severe damage resulting in enormous loss of several million pounds annually in the Egyptian economy. However, different controlling methods were used to bride the magnificent hazardous impacts of this pest. Such methods were carried out mechanically, chemically and by ionizing radiation and radioisotopes but the chemical controlling method by insecticides was the dominant currently used method. From the data presented in Table 7, it might be concluded that, all the tested complexes induced larval mortality of the cotton leaf worm, S. littoralis at all the tested concentration. All the tested compounds elicited 100% larval mortality at 100 ppm also, the complexes 10, 4, 9, 12 and 24 induced 100% larval mortality at 75 ppm. The complex no 4 gave the best result, which induced 77, 90, 100 and 100 larval mortality at 25, 50, 75 and 100 ppm, respectively. The rest complexes 7, 10, 12, 8 and 24 gave nearly similar effects. Such increased activity of these complexes could be explained on the basis of overtone’s concept (Anjaneyulu and Rao, 1986) and the tweedy’s chelation theory (Mishra and Singh, 1993). According to overtone’s concept of cell permeability, the lipid membrane that surrounds the cell favors the passage of only the lipid-soluble materials due to liposolubility which is an important factor, which controls the activity. On chelation, the polarity of the metal ion reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups. On the other hand, it increased the delocalisation of π-electrons over the whole chelate ring and enhanced the lipophilicity of the complexes. Therefore, the penetration of the complexes into lipid membranes was enhanced and the metal binding sites in the enzymes of the organisms were blocked. These complexes also disturbed the respiration process of the cell and thus block the synthesis of the proteins. Therefore, further growth of the organism was restricted. Furthermore, the mode of action of the complexes might involve the formation of a hydrogen bond through the functional groups of the complexes with the active centers of the cell constituents, resulting in interference with the normal cell process (Abdel-Hafew, 1978; Dharmaraj et al., 2001). In general, all tested complexes affected the survival of the cotton leaf worm, S. littoralis. In conclusion, the complexes 4, 7 and 10 proved to be promising controlling agents against this pest after testing their toxicity on mammals.

Table 7 Effect of the tested compounds on the larval mortality of Spodoptera littoralis.
Conc. 25 ppm 50 ppm 75 ppm 100 ppm
No.
Control
(2) 42 60 97 100
(4) 77 90 100 100
(7) 60 80 87 100
(8) 42 87 97 100
(9) 37 72 100 100
(10) 57 80 100 100
(12) 42 95 100 100

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