Synthesis, spectral characterization, molecular modeling and antimicrobial activity of new potentially N2O2 Schiff base complexes

Journal of Molecular Structure 1054–1055 (2013) 239–250

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Synthesis, spectral characterization, molecular modeling and antimicrobial activity of new potentially N2O2 Schiff base complexes Omima M.I. Adly ⇑, Ali Taha, Shery A. Fahmy Department of Chemistry, Faculty of Education, Ain Shams University, Roxy, Cairo 11711, Egypt

h i g h l i g h t s

g r a p h i c a l a b s t r a c t

 A novel tetradentate N2O2 Schiff base

ligand was prepared.  The Schiff base ligand was allowed to

react with a variety of metal ions.  Complexes have been characterized

by the analytical and spectral methods.  Molecular orbital calculations were performed.  Some complexes showed variable antimicrobial activities.

Molecular modeling of [Ni(L)].H2O

a r t i c l e

i n f o

Article history: Received 20 May 2013 Received in revised form 21 July 2013 Accepted 23 September 2013 Available online 2 October 2013 Keywords: N2O2 tetradentate Schiff base 5-Acetyl-4-hydroxy-2H-1,3-thiazine2,6(3H)-dione Molecular modeling Antimicrobial activity

a b s t r a c t Metal complexes of a new potentially tetradentate symmetrical Schiff base ligand (H2L) with Cu(II), Ni(II), Co(II), VO(IV), Zn(II), Cd(II), Ce(III), Fe(III) and UO2(VI) metal ions have been synthesized and characterized based on their elemental analyses, spectral (IR, UV–Vis, 1H NMR and mass spectra), magnetic and molar conductance studies as well as thermal gravimetric analysis (TGA). The synthesized complexes have the general formula [MHxL(H2O)yXn]: x = 0–1, y = 0–4 and n = 0–1; where: L = dianion of 6-hydroxy-5-[N-(2{[(1E)-1-(6-hydroxy-2,4-dioxo-3,4-dihydro-2H-1,3-thiazin-5-yl)ethylidene]amino}ethyl) ethanimidoyl]2H-1,3-thiazine-2,4(3H)-dione and X = nitrate or sulphate anion. The ligand behaves as diabasic tetradentate N2O2 sites, except in cases of Co(II), VO(IV) and UO2(VI) metal ions, it behaves as monobasic tetradentate Schiff base ligand. The metal complexes exhibited square planar, square-pyramidal and octahedral geometrical arrangements except for Ce(III) and UO2(VI) complexes, they are octa-coordinated. The Coats-Redfern equation was used to calculate the kinetic and thermodynamic parameters for the different thermal decomposition stages of some complexes. Structural parameters of the ligand and its metal complexes have been theoretically computed on the basis of semiemperical PM3 level, and the results were correlated with their experimental data. The antimicrobial activities of the ligand and its metal complexes were tested against some Gram-positive and Gram-negative bacteria; and fungus strain and the results were discussed. Ó 2013 Elsevier B.V. All rights reserved.

1. Introduction Tetradentate Schiff bases with N2O2 donor atoms are well known to coordinate with various metal ions and have attracted ⇑ Corresponding author. Tel.: +20 1008693220. E-mail address: [email protected] (O.M.I. Adly). 0022-2860/$ - see front matter Ó 2013 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.molstruc.2013.09.037

a great deal of interest in recent years due to their rich coordination chemistry [1–8]. They are also used as catalysts in polymer and dyes industry, beside some uses as antifertility and enzymatic agents as well as designing metal complexes related to synthetic and natural oxygen carriers [9,10]. Schiff base metal complexes display wide range of biological applications such as anticancer, antibacterial, antivirus and antifungal agents [11–15].

240

O

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O

H N OH

S

+

H 2N

NH2

-H2O

OH HO

S

N

O

O

O

H N

H N

H3 C

H3C

O

2.4. Molecular modeling

S

An attempt to gain a better insight on the molecular structure of the ligand and its complexes, geometric optimization and conformation analysis have been performed using PM3 forcefield as implemented in hyperchem 7.5 [19]

O

N CH3

Scheme 1. Formation of the Schiff base, H2L, ligand.

2.5. Elemental analysis and physical measurements

In continuation of our earlier work on synthesis of metal complexes of the ligands derived from 5-acetyl-4-hydroxy-2H-1,3-thiazine-2,6(3H)-dione (AHDT) [16,17]. The present work report the synthesis of symmetrical tetradentate Schiff base (H2L) formed by the condensation of AHDT and ethylenediamine, (Scheme 1). The solid complexes of Cu(II), Ni(II), Co(II), VO(IV), Zn(II), Cd(II), Ce(III), Fe(III) and UO2(VI) ions with this ligand (H2L) have been prepared and characterized by different physico-chemical techniques. Molecular modeling carried out for the free ligand and its complexes and the theoretical results were correlated with the experimental data. The ligand and its metal complexes were screened for their antimicrobial activity against Staphylococcus aureus and Bacillus subtilis as Gram-positive bacteria, Escherichia coli and Salmonella typhimurium as Gram-negative bacteria, Candida albicans and Aspergillus fumigatus as fungal strains.

2. Experimental 2.1. Materials Carbon disulphide, malonic acid, potassium thioacynate and ethylenediamine were either Analar or Merck. Copper(II), nickel(II), cobalt(II), cadmium(II), ceruim(III), iron(III) and dioxouranium(VI) were used as nitrate salts and were Merck. Oxovanadium(IV) sulphate and zinc(II) acetate were BDH. Organic solvents (ethanol, methanol, diethylether, isopropanol, acetic anhydride, acetic acid, dimethylformamide (DMF) and dimethylsulfoxide (DMSO)) were reagent grade and used without further purification. 2.2. Synthesis of 6-hydroxy-5-[N-(2-{[1-(6-hydroxy-2,4-dioxo-3,4dihydro-2H-1,3-thiazin-5-yl)ethylidene]amino}ethyl)ethanimidoyl]2H-1,3-thiazine-2,4(3H)-dione 5-Acetyl-4-hydroxy-2H-1,3-thiazine-2,6(3H)-dione (AHTD) was prepared as described elsewhere [18]. A mixture of AHTD (3.74 g, 20.0 mmol) and ethylenediamine (0.60 g, 10.0 mmol), in isopropanol (30 mL) containing few drops of acetic acid, was heated under reflux for 30 min. Yellow crystals were obtained after cooling at room temperature. The product was filtered off and recrystallized from ethanol, yield 2.98 g (75%) and m.p. 197 °C. Scheme 1 illustrates the preparation of the H2L ligand. 2.3. Synthesis of the metal complexes To a methanolic solution of the appropriate metal salt (3 mmol, 20 mL), a mixture of lithium hydroxide (0.25 g, 6 mmol) and H2L ligand (1.19 g, 3 mmol), in molar ratio 2:1 (LiOH:H2L) in methanolic solution (20 mL) was added. The reaction mixture was refluxed for 3 h, except in cases of VO(IV) was stirred at room temperature for 3 h. The Cu(II) complex 1 prepared in the absence of LiOH and was refluxed for 2 h. The resulting precipitates were filtered off, washed with bidistilled water, methanol then ether and finally air-dried. Trials to prepare Cr(III) complex of H2L, was unsuccessful.

Elemental analyses of (C, H, N, S) were carried out at the Ministry of Defense, Chemical War Department. Analyses of the metal ions were carried out complexometrically [20,21]. FT-IR spectra (4000–400 cm1) were recorded as KBr discs using FTIR Nicolet IS10 spectrophotometer (cm1), of the ligand and its complexes. 1 H NMR (300 MHz) spectra were recorded using a Mercury300BB. Dimethylsulphoxide, DMSO-d6, was used as a solvent and tetramethylsilane (TMS) as an internal reference. The electronic spectra of the ligand and its metal complexes were carried out on a JASCO model V-550 UV–VIS spectrophotometer in the range 200–900 nm. ESR spectra were recorded on the Bruker, Model: EMX, X-band spectrometer. Mass spectra were obtained using GC-2010 Shimadzu Gas chromatography instrument mass spectrometer (70 eV). Magnetic susceptibilities of the complexes were measured by the Gouy method at room temperature using Johnson Matthey, Alfa product, Model No. (MKI), the diamagnetic corrections were calculated from Pascal,s constants for all atoms in the compounds [22]. Thermal gravimetric analysis (TGA) data was measured from room temperature up to 800 °C (10 °C/min). The data were obtained using a Shimadzu TGA-50H instrument. Melting points were measured on a Stuart SMP3 melting point apparatus. 2.6. Biological activity The standardized disc agar diffusion method [23] was followed to determine the activity of the synthesized compounds against the sensitive organisms Staphylococcus aureus and Bacillus subtilis as Gram-positive bacteria, Escherichia coli and Salmonella typhimurium as Gram – negative bacteria, Candida albicans and Aspergillus fumigatus as fungus strain. The compounds were dissolved in DMSO which has no inhibition activity to get concentration of 100 lg mL1. The test was performed on medium potato dextrose agars (PDA) which contain infusion of 200 g potatoes, 6 g dextrose and 15 g agar [24]. Uniform size filter paper disks (3 disks per compound) were impregnated by equal volume (10 lL) from the specific concentration of dissolved tested compounds and carefully placed on inoculated agar surface. After incubation for 36 h at 27 °C in the case of bacteria and for 48 h at 24 °C in the case of fungi, inhibition of the organisms was measured and used to calculate mean of inhibition zones. 3. Results and discussion All complexes were stable at room temperature and were soluble in DMF and DMSO but sparingly soluble in alcohols. The analytical and physical data of the free ligand, (H2L), and its solid complexes with Cu(II), Ni(II), Co(II), VO(IV), Zn(II), Cd(II), Ce(III), Fe(III) and UO2(VI) ions are presented in Table 1. The ligand (H2L) behaves either as monobasic or dibasic tetradentate via N2O2 sites (Scheme 1). 3.1. Mass spectral studies The mass spectrum of the Schiff base (H2L) ligand showed the molecular ion peak at m/e 398 which agrees well with the

S

— 13.34 (13.29) 12.86 (12.81) 12.99 (12.40) 10.24 (10.13) NA (8.06) 9.83 (9.86) NA (20.40) 17.25 (17.21) 12.54 (12.67) NA (30.73) 15.89 (16.09) 13.17 (13.41) 12.14 (12.92) 13.14 (13.55) 11.31 (11.04) 14.95 (15.20) 11.32 (11.32) 9.10 (9.34) 9.90 (9.82) 12.25 (12.43) 8.42 (8.28) 3.55 2.29 3.35 2.68 3.21 3.86 3.58 3.49 4.38 3.21 2.24 41.98 35.26 33.72 34.98 29.19 26.45 30.45 25.02 25.27 32.68 21.68 197 277 274 >300 >300 >300 >300 >300 >300 >300 >300 Pale yellow Violet Brown Red Brown Green Brown Yellow Yellow Yellow Yellow (75) (57) (68) (52) (64) (76) (59) (78) (63) (72) (71) NA: not analyzed.

(1) (2) (3) (4) (5) (6) (7) (8) (9) (10)

C14H14N4O6S2 C14H14N4O7 S2 Cu C14H16N4O8 S2 Cu C14H14N4O7S2Ni C14H20N5O12.5S2Co C14H21N4O15S3V C14.5H18N5O11.5S2Fe C14.5H22N5O13.5S2Ce C14H28N4O14S2Cd C14H18N4O9S2Zn C14H18N5O13.5S2U H2L [Cu(L)H2O] [Cu(L)(H2O)2] [Ni(L)]H2O [Co(HL)(H2O)(NO3)]2.5H2O [VO(HL)(SO4)]4H2O [Fe(L)(H2O)(NO3)]0.5CH3OHH2O [Ce(L)(H2O)4]NO30.5CH3OH [Cd(L)(H2O)2]6H2O [Zn(L)(H2O)2]H2O [UO2(HL)(H2O)(NO3)]1.5H2O

C

(42.20) (35.18) (33.90) (35.54) (28.94) (26.59) (30.75) (25.36) (25.75) (32.59) (21.71)

H

(3.54) (2.95) (3.25) (2.98) (3.46) (3.34) (3.20) (3.22) (4.32) (3.51) (2.34)

N

3.2. Infrared spectra

398.40 477.94 495.95 473.11 581.37 632.45 566.27 686.59 652.89 515.82 774.45

Elemental analyses found (calcd.) % M.P. (°C) Color M Wt. Yield (%) Molecular formula Ligand/complex

Table 1 Physical and analytical data of H2L ligand and its metal complexes.

241

molecular formula C14H14N4O6S2. The fragmentation pattern of the mass spectrum is depicted in Scheme 2.

13.83 (14.06) 11.71 (11.72) 11.62 (11.29) 11.52 (11.84) 11.90 (12.05) 8.81 (8.85) 12.69 (12.36) 10.43 (10.20) 8.81 (8.58) 10.92 (10.86) 9.13 (9.04)

M

O.M.I. Adly et al. / Journal of Molecular Structure 1054–1055 (2013) 239–250

The significant IR spectral data of the free ligand (H2L) and its metal complexes and their tentative assignments are listed in Table 2. Comparing the IR spectrum of the Schiff base ligand with both that of 5-acetyl-4-hydroxy-2H-1,3-thiazine-2,6(3H)-dione (AHTD) and ethylenediamine, was important to assign the IR absorption bands of the current ligand. The intense band at 1612 cm1 was assigned to the azomethine m(C@N) mode. This assignment further supported by the disappearance of the absorption band that assigned to the acetyl C@O of (AHTD) at 1659 cm1, which is condensed with the NH2 group of the ethylenediamine as suggested in Scheme 1 [17,25]. The free ligand exhibits two partially resolved bands at 1685 and 1625 cm1 which assigned to the stretching frequencies of the two carbonyl groups in the thiazine ring (mC@O), the former arises from the carbonyl adjacent to carbon and the second adjacent to NH (as tautomers) with splitting Dm = 60 cm1. Comparison of the IR spectra of the metal complexes with that of the free ligand (H2L) revealed that all complexes showed broad diffuse band in the 3366–3524 cm1 region which may be assigned to mOH of the coordinated/solvated water or methanol molecules and/or the coordinated phenolic oxygen. It is difficult to distinguish between the stretching vibrational m(OH) of phenolic groups of the primary ligand and that of water or methanol molecules because the overlapping values and located in one place. To ascertain the involvement of phenolic group in the coordination process, this requests to make a follow up of the stretching vibration bands of m(CAO) in all complexes; this interpretation further confirmed by the appearance of m(MAO) in the region (473–410) cm1 [26]. This check shows that the m(CAO) was shifted from 1067 cm1 of the free ligand to (1072–1095) cm1 in the complexes along with the decreasing of their intensities, pointing to the involvement of the phenolic oxygen in the coordination sphere [27,28]. This emphasized from the linear relationship between m(MAO) and m(CAO) phenolic as shown in Fig. 1, m(CAO) phenolic/cm1 = 1326.4– 0.539 m(MAO)/cm1, R2 = 0.94, (n = 6 points except 3, 4, 7 and 10 complexes). The negative slope reveals decreasing of (CAO) phenolic bond strength with the increasing of interaction between metal ion and phenolic oxygen, this agree well with the Gutmann’s bonds rule [29]. The intense band at 1612 cm1 that assigned to the azomethine (C@N) group in the free ligand was shifted to lower frequencies in the range (11–53 cm1), indicating the involvement of azomethine (C@N) group in the coordination sphere [30]. This emphasized by the appearance of a new weak to medium intensity absorption bands in the region (547–510) cm1 that may be attributed to m(MAN) [31,32]. The linear correlation of m(C@N) versus m(CAO) phenolic, m(C@N)/cm1 = 5400–3.54 m(CAO))/cm1 phenolic, R2 = 0.91, (n = 5 points except 4, 6, 7, 9 and 10 complexes), the negative slope, indicates increasing the bond strength of the azomethine (C@N) group accompanied by decreasing the bond strength (CAO) phenolic. Moreover, the positive slope of m(MAN) versus m(MAO), m(MAN) = 13.52 + 1.19 m(MAO), R2 = 0.97, (n = 6 points except 3, 7, 8 and 10 complexes), further support for this point of view as shown in Scheme 3. This point of view further evidenced by the negative slope of the relationship of mC@N versus DmC@O (splitting of carbonyl stretching) of complexes yielding; mC@N/cm1 = 1622-1.149 DmC@O/cm1, R2 = 0.99 (n = 5 points except 2, 3, 6, 7, 10); revealing increase the extent of interaction between the metal ion and azomethine nitrogen leads to decrease the splitting of stretching frequencies of the carbonyl groups (DmC@O).

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N H

b

S

HO

OH

a

398 (14%)

O H3C

O

b N

N

S O

+.

CH3

O H3C

N H

S

+

O

-S

+.

O

O

70 (17)

CH3

O H3C

N+ O

N H

O

102 (3)

N

N

S

+

S

-32

HO

OH 295 (15)

-16 -0.5 O2

O

N

N S

+ O

CH3

H3C

N

N S

+ O

S N H

O

235 (18) -28

H

OH 279 (8)

-CO

.

-C2H, -CH3

-40 CH3

H3C

+.

N

N

.

-44

-24 -C2

O

-NH, -CO, -H

+.

CH3

O H3C

CH3

H3C

+

N OH H 185 (14)

O

O

a N

N

S O

O H3C

N

N

.

+

N H

O 207 (15)

211 (14)

C

N

N

S

+

OH 239 (17)

-16 -0.5 O2 CH3

H3C S

N

N

H

N

N

H3C

C

+

96 (100)

+ 191 (4)

Scheme 2. Mass fragmentation pattern of the Schiff base, H2L, ligand.

Table 2 Characteristic infrared frequencies (cm1) of Schiff base, H2L, ligand and its metal complexes. Ligand/complex

m(OH, NH)

m(CAH)

m(C@O)

m(C@N) m(CAO) m(MAN) m(MAO) Other band

aliphatic (1) (2) (3) (4) (5)

H2L [Cu(L)H2O] [Cu(L)(H2O)2] [Ni(L)]H2O [Co(HL)(H2O)(NO3)] 2.5H2O [VO(HL)(SO4)]4H2O

3420 br, 3278 br 3524,3275 br, 3092 br 3504 br, 3153 br 3446 br, 3150 br 3420 br, 3245 br 3446 br, 3167 br

2962 2833 2955 2823 2955 2955

w w w w w w

1685 1682 1667 1683 1667 1652

s s s s s s

(6) (7) (8) (9) (10)

[Fe(L)(H2O) (NO3)]0.5 CH3OH.H2O [Ce(L) (H2O)4]NO3. 0.5CH3OH [Cd(L)(H2O)2]6H2O [Zn(L)(H2O)2]H2O [UO2(HL)(H2O)(NO3)]1.5H2O

3435 3366 3393 3420 3470

2955 2955 2924 2940 2970

w w w w w

1685 1670 1667 1675 1683

s 1645 s s 1619 s br 1650 s s 1636 s s 1615 s

br, 3277 br br br br br

1625 1656 1632 1642 1621 1635

s s s s s s

1612 s 1595 m 1598 s 1559 s 1570 s 1601 s

1067 1073 1078 1084 1057 1072

w w w w w w

— 544 536 543 530 545

w w w w w

— 466 465 419 455 467

w w w w w

1575 s 1598 m 1559 s 1576 s 1574 s

1072 1095 1083 1087 1072

w w w w w

545 530 543 510 547

w w w w w

473 467 455 442 460

w w w w w

— — — — 1344, 1215 NO3 (unidentate) 970 w (V@O) 1114 w, 1030 w m 3 SO4 (unidentate) 1350, 1270 NO3 (unidentate) 1384, 835 NO3 (ionic) — — 1351, 1270 NO3 (unidentate) 915 w (O@U@O)

s = strong, m = medium, w = weak, br = broad.

The anion coordination modes were inferred from the IR data of the current complexes. Nitrate anion in complexes 4, 6 and 10 showed two absorption bands at (1344, 1215), (1350, 1270) and (1351, 1270) cm1, respectively, indicating that the nitrate behaves as unidentate nature. However, in complex 7 the NO3 group acts as ionic group because it displays two bands at (1384 and 835) cm1.

In complex 5, when the sulphate anion bound to the metal symmetry of sulphate anion changes from Td to C3v. The attached metal shifts the unique SAOAM stretch to low frequency. The three remaining SAO bonds experience a C3v pseudo symmetry and they behave like a1 + e stretches. Both are IR allowed, so there should be two vibrations in the SAO stretching region (1030–1114) cm1

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970 cm1 due to the stretching vibration of the m(V@O) band in hexa-coordinated structure [33]. Similarly, the strong band located at 915 cm1 for complex 10 was assigned to m3(asym) of the O@U@O group [34,35].

uM-O/cm-1

475

465

455

3.3. 1H-NMR spectra

445 C-O

435 1070

1075

1080

phenolic

1085

1090

Fig. 1. Relationship of m(MAO) bond versus m(CAO) phenolic frequencies.

O

O HN

NH O

S

O

S

M O

N

O

N

1 H-NMR spectral data of the organic ligand (H2L) and its Ni(II) and Zn(II) complexes are depicted in Table 3. The methyl protons were appeared at d 2.51, 2.25 and 2.31 ppm in the 1H-NMR spectra of the ligand and its Ni(II) and Zn(II) complexes, respectively. While, the methylene protons were appeared at d 3.88, 3.58 and 3.75 ppm, respectively. Also, the NHthiazine protons appeared as a broad exchangeable signal at d 11.64, 11.80 and 11.20 ppm in the ligand and its Ni(II) and Zn(II) complexes, respectively. On the other hand, the OHthiazine protons appeared as a broad exchangeable signal at d 12.52 ppm in the spectrum of free ligand which were disappeared in the spectra of the Ni(II) and Zn(II) complexes. These results confirm the deprotonation of the phenolic group during the coordination of Ni(II) and Zn(II) ions with the free ligand.

CH3

CH3

3.4. Electronic spectra and magnetic studies

Decrease bond force Increase bond force Scheme 3. Effect of coordination to metal ion on the bond force.

Table 3 H NMR spectra of the Schiff base H2L ligand and its Ni(II) and Zn(II) complexes (in DMSO-d6). 1

1

H NMR Chemical shifts in ppm

H2L ligand

Ni(II) complex

Zn(II) complex

Assignment

2.51 3.88 11.64 12.52

2.25 3.58 11.80

2.31 3.75 11.20

(s, 6H, 2CH3) (s, 4H, 2CH2) (bs, 2H, NHthiazine exchangeable with D2O) (bs, 2H, OHthiazine exchangeable with D2O)

emphasizing the unidentate nature of sulphate anion [31,33]. The IR spectrum of VO(IV) complex 5 displayed absorption band at

The electronic spectral data of the free ligand and its metal complexes, in the range 200–900 nm, in DMF solvent and Nujol mull as well as their magnetic susceptibility (leff) and molar conductance values are displayed in Table 4. Electronic spectrum of the free ligand in DMF solution displayed two absorption bands at 35,211 and 30,864 cm1, the first band could be attributed to the p ? p transition of the azomethine group (K-band), and the second band might arises from the n ? p transitions resulting from nitrogen, oxygen and sulfur atoms (R-band) [34]. These bands exhibit more or less shift in complexes. A comparison of the electronic spectra of the free ligand with those of their corresponding metal complexes show some shifts which can be considered as evidence for the complex formation. The electronic spectra of Cu(II) complexes 1 and 2 exhibit two bands as; an asymmetric broad bands at 17,361 and 17,453 cm1, respectively, and a more intense bands at 24,390 and 24,510 cm1. The latter band might assign to ligand–metal charge transfer transition. The asymmetric band is assigned to 2 2 d–d (2A1 B1) and/or (2Eg T2g) transition. The band position

Table 4 Characteristic electronic transition bands in DMF and magnetic moments of the metal complexes of H2L ligand. Ligand/complex H2L (1)

[Cu(L)H2O]

(2)

[Cu(L)(H2O)2]

(3)

[Ni(L)] H2O

(4)

[Co(HL)(H2O)(NO3)] 2.5H2O

(5)

[VO(HL)(SO4)]4H2O

(6)

[Fe(L)(H2O)(NO3)]0.5 CH3OH.H2O

(7)

[Ce(L)(H2O)4] NO3 0.5CH3OH

(8)

[Cd(L)(H2O)2] 6H2O

(9)

[Zn(L)(H2O)2] H2O

(10)

[UO2(HL)(H2O)(NO3)] 1.5H2O

DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol DMF Nujol

p–p (Nujol)

n  p

d–d Transition bands

K Ohm1 cm2 mol1

leff

35,211 (36,232) — (36,231) —

30,864 (31,746) 32,468 (32,573) 32,467 (32,679) 28,571 (29,155) 29,411 (29,155) 31,250 (32,051) — (28,985) 27,397 (27,322) 27,174 (27,247) 27,778 (27,700) 29,411 (30,211)

—

—

—

24,390, 17,361 (26,809, 18,051) 24,509, 17,452 (25,839, 17,668) 21,692, 17,361 (22,026, 17,301) 23,980, 18,518 (23,364,18,281) 23,529, 19,607 (23,474, 19,011) 20,202, 17,857 (20,040, 17,793) 23,809 — —

3.24

1.13

5.33

2.12

4.17

Dima.

20.02

4.70

12.68

1.33

36.00

3.02

60.68

2.37

11.18

—

—

5.60

—

22,573 (22,727)

4.64

—

— — (33,670) 35,460 33,444 (34,246) 35,714 (35,587) 34,722 (34,843) 34,722 (34,722) 33,898 (35,087)

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ecule in addition to one water molecule and one nitrate anion. Furthermore, the mass spectrum of complex 10 (Supplementary Fig. IIS) showed the parent peak at m/e 774 which compares well with the suggested formula weight of the complex. 3.5. Conductance studies The molar conductance values, in DMF (103), of all complexes are shown in Table 4 indicating that all complexes are neutral in the range (3.24–36.00) except complex 7 has molar conductance values 60.70 which suggest the ionic nature of this complexes (1:1 electrolytes) [49]. 3.6. Thermal gravimetric analysis From the TG, and DTG curves of the ligand and its complexes some information and physical parameters are obtained, such as: thermal stability of ligand and its complexes, the presence or absence of coordinated water molecules, thermal decomposition. The decomposition stages, determined temperature ranges and the corresponding percent mass losses (experimental and calculated ones) are given in (Supplementary Table 1S). Figs. 2 and 3 represent complexes (1, 3, 6, 7); [Cu(L)H2O], [Ni(L)].H2O, [Fe(L)(H2 O)(NO3)]0.5CH3OH. H2O and [Ce(L)(H2O)4]NO3.0.5CH3OH, respectively. The results reveal good agreements with the proposed formula weight expected from the analytical and spectral data. Thermal gravimetric analysis (TGA) of the free ligand H2L undergoes three stages of decomposition in the range 114–641 o C. The first stage in the temperature range 114–232 °C corresponding to the mass loss (calcd./found%; 27.61/27.50%) may be attributed to the decomposition of two molecules of CH3CN and one molecule of C2H4 at DTmax = 197 °C. The second stage within the temperature range 232–392 °C with mass loss (calcd./found%; 35.64/35.34%) indicating the removal of four molecules of CO and one molecule of N2 and H2 at DTmax = 338 °C. The third stage observed within the range 392–641 °C with mass loss (calcd./found%; 24.59/25.00%) matching with one molecule of S2, O2 and H2 at DTmax = 549 °C, leaving 4C as deposit. The Cu(II) and Ni(II) complexes 1 and 3 thermograms showed three stages of decomposition pattern as shown in (Fig. 2A and B). The first stage within the temperature range 63–217 °C and 37–156 °C with mass loss (calcd./found%; 3.76/3.40%) and (calcd./ found%; 3.80/3.82%), respectively, indicating the removal of one molecule of coordinated water and one molecule of lattice water 100

TGA (mg)

and the magnetic moment values (1.13 and 2.12 BM) can be taken as an evidence for the square-pyramidal and/or octahedral geometries [36,37]. The first value (1.13 BM) was found much lower than expected for d9 system, indicating to a strong metal–metal interaction between the complex molecules, whereas the second one (2.12 BM) is little higher than expected [38]. The electronic spectrum of nickel(II) complex 3, showed a sig1 nificance at 17,361 cm1 which is assigned to 1A2g A1g (i.e. dx2y2 dxy) transition in square planar structure [39,40]. The magnetic susceptibility of the complex indicate that leff = 0.00 BM, so the complex is diamagnetic. It is considered as an additional evidence for square planar geometry around Ni(II). The other bands appeared at 28,571 and 21,392 cm1 have been assigned to LMCT and MLCT, respectively [41]. The electronic spectrum of Co(II) complex 4 exhibits two bands 4 at 23,980 and 18,519 cm1 assigned to 4T1g(P) T1g(F) (m3) and 4 4 A2g(F) T1g(F) (m2) transitions around the Co(II) in octahedral geometry [38]. The m1 band not observed, but it can be calculated and found at 13,193 cm1. The ligand field parameters were calculated B (420 cm1), 10Dq (1176 cm1) and b (0.43) and found within the range reported for octahedral structure. The leff value (4.70 B.M.) and color of Co(II) complex 4 are further support the octahedral geometry around cobalt(II) ion [42,43]. The green oxovanadium(IV) complex 5 showed a well defined 2 band at 19,607 cm1 assigned for 2B2 B1 electronic transition, suggesting hexa-coordinated structure, which would be octahedral [44]. This conclusion further confirmed by the IR spectrum of vanadyl complex 5 that exhibited a band at 970 cm1 as shown in Table 2, which is characteristic for hexa-coordinated geometry for vanadyl complexes [45]. The subnormal magnetic moment, leff = 1.33 B.M, may propose a large metal–metal interaction between the complex molecules such that observed for the present Cu(II) complex 1. The electronic spectrum of Fe(III) complex 6 is dominated by two absorption bands observed at 20,202 and 17,857 cm1. The first band is assigned to a charge transfer band (MLCT) and the second band could be assigned to d ? d (6A1(G) ? 4T2(S)) electronic transition provides evidence for octahedral arrangement around the Fe(III) ion, emRax = 113 and 57 L mol1 cm1, respectively [43]. The value of the magnetic moment of Fe(III)-complex 6 is 3.02 B.M (Table 4), suggests high spin M low spin equilibrium (t2g3eg2 M t2g5eg0) of octahedral species in an approximate ratio 1:1 [42]. The mass spectrum of the complex 6 (Supplementary Fig. IS) showed the parent peak at m/e 566 which compares well with the formula weight of the complex. The electronic spectrum of the Ce(III) complex 7 showed a new absorption band at 23,810 cm1 which may be related to MLCT transition [46]. The magnetic moment of the complex is 2.37 B.M. which is close to the experimental range of 2.14–2.46 B.M. [47,48]. The electronic spectra of yellow Cd(II) and Zn(II) complexes 8 and 9 reveal a charge transfer bands p ? p and n ? p transitions in the vicinity of the Schiff base ligand at 34,722 for both complexes; and 27,778 and 27,174 cm1, respectively suggesting an octahedral structure which is common for d10 systems. The electronic spectrum of the orange red UO2(VI) complex 10 exhibits three bands observed at 33,898, 29,412 and 22,573 cm1. The former two bands coincide with the bands observed in the free ligand with a slightly red shift because of the coordination with the UO2(VI) ion. The later broad less intense band observed at 22,573 cm1 can be assigned to metal ? ligand charge transfer and/or apical oxygen ? fo (U) transition which is allowed transition producing the orange red color [45]. The coordination sphere of UO2(VI) can be described as octa-coordinated with the two oxygen in the apical positions with a planar equatorial environment containing a mononegative tetradentate ligand mol-

80 60 40

(B)

20 0

100

200

300

400

500

600

700

800

900

Temp (C) 2.0

TGA (mg)

244

1.5 1.0 0.5

(A)

0.0 0

100

200

300

400

500

600

700

800

Temp (C) Fig. 2. TGA curves of [Cu(L)(H2O)] (A) (1) and [Ni(L)]H2O (B) (3).

900

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tively. The third step found in the temperature ranges 165–287, 264–475 and 149–274 °C analogous to mass loss (calcd./found%; 17.27/18.79%), (calcd./found%; 21.32/20.62%) and (calcd./found%; 13.17 13.58%) representing removal of (one molecule of coordinated water and two molecules of CH3CN), (two molecules of CH3CN and one molecule of C2H4) and (one molecule of coordinated water two CH3CN and one H2 molecules) at DTmax = 246, 387 and 242 °C for complexes 4, 9 and 10, respectively. The fourth step assigned within the temperature ranges 287–404, 475–690 and 274– 501 °C equivalent to mass loss of (calcd./found%; 16.02/16.20%), (calcd./found%; 21.32/20.62%) and (calcd./found%; 26.21/27.08%), representing removal of (one of HNO3, one N2 and one H2 molecules), (four CO gas molecules); and (one molecule of HNO3, four CO and one N2 molecules) at DTmax = 348, 617 and 386 °C, leaving behind C10H4O6S2Co, C4H2O2N2S2Zn and C6H4O3S2U species for complexes 4, 9 and 10, respectively. The TG of Fe(III) and Ce(III) complexes 6 and 7 display three stages. The first stage observed at temperature ranges: 33–80 and 33–119 °C with mass loss (calcd./found%; 2.82/2.64%) and (calcd./found%; 2.32 /2.55%) indicating the removal of half molecule of lattice methanol at DTmax = 55 and 54 °C for both complexes 6 and 7, respectively. The second step established at temperature ranges: 83–273 and 123–338 °C related to mass loss (calcd./ found%; 30.02/29.89%) and (calcd./found%; 22.42/22.40%) reveal elimination of (one lattice water, one coordinated water, two CH3CN and C4H4 molecules) and (four coordinated water and two CH3CN molecules) at DTmax = 228 and 272 °C for complexes 6 and 7, respectively. The third step appears at temperature ranges: 274–511 and 339–693 °C with mass loss (calcd./found%; 33.39/ 34.18%) and (calcd./found%; 13.25/14.00%) which corresponds to decomposition of the most part of the ligand (one of HNO3, four CO and half N2 molecules) and (one of HNO3 and one N2 molecules) at DTmax = 364 and 473 °C, leaving at the back C2O2NS2Fe and C10H5O6S2Ce species for complexes 6 and 7, respectively. The kinetic and thermodynamic parameters viz order of reaction (n), energy of activation (Ea), frequency factor (log A), entropy of activation (DS), enthalpy change of activation (DH), and free energy change of activation (DG) for non-isothermal decomposition of metal complexes have been determined for the first two stages by Coats-Redfern integral method [50] and the data is given in Table 5. The results showed that the calculated free energy of activation (DG) is relatively low and positive sign indicating the autocatalytic effect of metal ions on thermal decomposition of the complexes and non-spontaneous processes [51]. DS values were negative, which indicate a more ordered activated state that may be possible through the chemisorptions of oxygen and other decomposition products. The more ordered nature may be due to the polarization of bonds in activated state which might happen

TGA (mg)

1.2 1.0 0.8 0.6 0.4

(B)

0.2 0.0 0

100

200

300

400

500

600

700

800

900

Temp (C)

TGA (mg)

2.0 1.5 1.0 0.5

(A)

0.0 0

100

200

300

400

500

600

700

800

900

Temp (C) Fig. 3. TGA curves of [Fe(L)(H2O)(NO3)]0.5CH3OH.H2O (A) (6) and [Ce(L)(H2O)4]NO30.5CH3OH (B) (7).

at DTmax = 194 and 113 °C, respectively. The second stage within the temperature ranges 217–263 °C and 157–292 °C, analogues to the loss (two molecules of CH3CN, one ethane (C2H4), N2 and H2 gas molecule) and (two molecules of CH3CN and C4H4 species) corresponding to (calcd./found%; 29.29/31.00%) and (calcd./ found%; 28.32/27.55%), at DTmax = 242 and 263 °C, respectively. The third stage observed in the temperature ranges 263–422 °C and 292–532 °C, corresponding to (four molecules of CO and one molecule of N2 and H2) and (four molecules of CO and one molecule of N2, S2 and H2) leaving C4O2S2Cu and C2O2Ni species for Cu(II) and Ni(II) complexes 1 and 3 at DTmax = 391 and 494 °C, respectively. The thermograms of Co(II), Zn(II) and UO2(VI) complexes 4, 9 and 10 exhibit four distinct steps. The first step observed at temperature ranges: 33–105, 99–178 and 58–92 °C equivalent to mass loss (calcd./found%; 4.64/4.44%), (calcd./found%; 3.48/3.13%) and (calcd./found%; 1.16/1.07%) revealing the removal of 1.5 molecules, one molecule and 0.5 molecule of lattice water at DTmax = 71, 113 and 75 °C for complexes 4, 9 and 10, respectively. The second step established at the temperature range 105–164, 178–264 and 93– 148 °C corresponding to mass loss (calcd./found%; 3.10/2.57%), (calcd./found%; 6.97 /6.87) and (calcd./found%; 2.32/2.00%) indicating elimination of one molecule of lattice water, two molecules of coordinated water molecules and one molecule of lattice water at DTmax = 144, 251 and 123 °C for complexes 4, 9 and 10, respec-

Table 5 Temperatures of decomposition and the kinetic parameters of complexes. Compound [Cu(L)(H2O)] (1) [Ni(L)].H2O (3) [Co(HL) (H2O) (NO3)] 2.5H2O (4)

Fe(L)(H2O) (NO3)].0.5 CH3OH.H2O (6) [Ce(L)(H2O)4].NO3.0.5CH3OH (7) [Zn(L)(H2O)2].H2O (8) [UO2(L) (H2O) (NO3)].1.5H2O (10)

Step 1st 2nd 1st 2rd 1st 2nd 3rd 1st 2nd 1st 1st 2rd 1st 2rd

n Order 0.66 0.66 1 0 1 0 0 1 0.66 0.5 0.5 0 0.66 0

T (K) 467 515 386 536 56.45 417 519 328 501 327 386 624 348 396

A (S1) 8

2.2  10 4.1  1013 4.4  108 1.2  106 1.2  108 1.9  108 2.64  107 1.9  109 2.5  105 1.1  109 2.1  107 1.3  108 2.30  109 3.63  107

DE(kJ mol1)

DH (kJ mol1)

DS (kJ mol1 K1)

DG (kJ mol1)

14.11 137.9 68.96 52.71 56.45 5.98 26.93 59.20 45.97 3.74 18.29 16.54 64.68 15.63

10.22 133.64 65.75 48.25 53.59 2.51 22.61 56.47 41.80 1.02 15.08 11.35 61.78 12.33

0.097 0.004 0.088 0.141 0.099 0.097 0.115 0.113 0.154 0.080 0.115 0.103 0.075 0.11

55.52 131.51 99.72 124.27 87.64 42.95 82.29 93.53 118.95 27.18 59.47 75.62 87.88 56.25

2400

2600

2800

3000

3200

3400

3600

3800

4000

4200

Field (G) Fig. 4. ESR spectra of [Cu(L)(H2O)2] (A) (2) and [VO (HL)(SO4)]4H2O (B) (5).

– – – 16 I 8L 6L 21H 6L – 15 I 14 I 27 – – – 16 I 12 L 8L 27H 7L – 20 I 18I 38 – – – 4L 13 I 7L 7L 9L 8L 18 I 16 I 28 – – – 6L 15 I 8L 10 L 16 I 9L 22 I 22 I 36 – – 12 I 6L 14 I 4L 6L 5L 9I 18 H 9L 25 – – 14 I 6L 16 I 7L 9L 6L 15 I 25H 6L 35 – – 12I – 5L – 18 H 9L – 4L 7L 26

0.5 mg/ml 1 mg/ml 0.5 mg/ml 1 mg/ml 0.5 mg/ml 0.5 mg/ml

– = No effect. L: Low activity = Mean of zone diameter 61/3 of mean zone diameter of control. I: Intermediate activity = Mean of zone diameter 62/3 of mean zone diameter of control. H: High activity = Mean of zone diameter >2/3 of mean zone diameter of control. #: Chloramphencol in the case of Gram – positive bacteria, Cephalothin in the case of Gram – negative bacteria and cycloheximide in the case of fungi. Calculate from 3 values. ** Identified on the basis of routine cultural, morphological and microscopical characteristics.

(A)

– – 20I – 6L – 25H 5L – 10 L 8L 35

4200

H2L [Cu(L)H2O] [Cu(L)(H2O)2] [Ni(L)]H2O [Co(HL)(H2O)(NO3)]2.5H2O [VO (HL)(SO4)]4H2O [Fe(L)(H2O)(NO3)]0.5 CH3OH.H2O [Ce(L) (H2O)4]NO3. 0.5CH3OH [Cd(L)(H2O)2]6H2O [Zn(L) (H2O)2]H2O [UO2(HL) (H2O) (NO3)]1.5H2O Control #

4000

1 2 3 4 5 6 7 8 9 10

3800

1 mg/ml

3600

1 mg/ml

3400

Field (G)

Salmonella typhimurium (ATCC 14028)

3200

Bacillus subtilis (ATCC 6635)

3000

Staphylococcus aureus (ATCC 25923)

2800

Gram – positive bacteria

2600

Table 6 The antimicrobial activity of the newly synthesized compounds.

2400

Organism

(B)

Mean* of zone diameter, nearest whole mm

The X-Band ESR spectra of [CuL(H2O)2] (2) and [VO(HL)(SO4)].4H2O (5) complexes at room temperature were shown in (Fig. 4). The spectrum of the Cu(II) complex 2 exhibits two broad band with g|| = 2.17 and g\ 2.07 so, g|| > g\ > 2.0023, indicating that the unpaired electron of Cu(II) ion is localized in the dx2–y2 orbital [52]. In axial symmetry, the g-values are related to the G-factor by the expression G = (g||  2)/(g\  2) = 4. According to Hathaway and Billing [53]. The G values of the complex 2 are <4 suggesting that the considerable exchange interaction in the solid state. Further, the shape of the ESR spectra of complex 2 indicates that the geometry around the Cu(II) ions are elongated octahedron [54,55]. Molecular orbital coefficients, a2 and b2 were calculated as described previously [30]. The lower value of a2 (0.44) compared to b2 (1.00) in complex 2 indicate that the covalent in-plane r – bond-

Gram - negative bacteria

Escherichia coli (ATCC 25922)

3.7. ESR spectral studies

*

– – 7L – – – – – – 18 H – 26 – – 9L – – – – – – 25H – 37 – 13 I 15 I – 9L 8L 17 I 14I 6L 22 H – 28 – 16 I 18 I – 15 I 12 I 18 I 14 I 7L 24 H – 35

1 mg/ml 0.5 mg/ml 1 mg/ml

Yeasts and Fungi**

Aspergillus fumigatus

through charge transfer electronic transitions. The positive values of DH means that the decomposition processes are endothermic. The linear relationships of the enthalpy change of activation (DH) versus the calculated heat of formation of the complexes (DHf/k cal/mol) or the energy gap (Egap = ELUMO–EHOMO), yield the following correlations: DHf/k cal mol1 = 13.48–2.38 DH/k cal mol1, R2 = 0.99 (n = 3 except complexes 1 and 3) and Egap/ k cal mol1 = 37.51 + 0.095 DH/k cal mol1, R2 = 0.97, (n = 4, except complex 8). The negative slope of the former relationship reveals decreasing the enthalpy change of activation (DH) with the increasing heat of formation. However, the positive slope of the later correlation indicates increasing the enthalpy change of activation with the increasing of Egap (measures the stabilization energy of the complex) vide infra.

0.5 mg/ml

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Candida albicans (ATCC 10231)

246

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ing is more pronounced than the covalent in-plane p-bonding character. The ESR spectrum of the oxovanadium(IV) complex 5, exhibits two bands at g\ = 2.07 and g|| = 2.16 indicating an octahedral geometry around VO(IV) ion confirming the results obtained from

the IR and UV/Vis spectra. The absence of vanadium’s hyperfine coupling which is common in the solid state might attribute to the strong exchange interactions, which average out the interaction with the nuclei or due to simultaneous flipping of neighboring electron spin [17,56].

Table 7 Molecular orbital parameters of the Schiff base H2L ligand and its metal complexes. Property

H2L

1

2

3

4

6

8

9

Total energy (kcal/mol) Heat of formation (kcal/mol) Dipole moment (D) HOMO (eV) LUMO (eV) Egap (eV)

107,378 147.23 6.34 9.75 1.69 8.06

141,565 313.49 8.64 8.18 1.24 6.94

149,056 365 9.89 8.36 1.20 7.16

130,789 365 8.93 9.11 1.31 7.8

156,515 574 9.21 9.72 1.86 7.86

150,637 630 8.90 9.37 1.65 7.72

122,153 207 3.29 9.22 1.13 8.09

122,280 215 6.05 9.35 1.42 7.93

Table 8 Theoretical calculated the bond lengths (A°) of the Schiff base ligand (H2L) and its metal complexes on the PM3 level.

(1) (2) (3) (4) (6) (8) (9)

Compound

C(24)@N(25) C(10)@N(12) C(2)@O(8) C(20)@O(21) O(23)@H(37) O(9)@H(28) MAO(9) MAO(23) MAN(12) MAN(25) MAOsolv Other

H2L [Cu(L) H2O] [Cu(L) (H2O)2] [Ni(L)].H2O [Co(HL) (H2O)(NO3)] 2.5H2O [Fe(L)(H2O)(NO3)]0.5 CH3OH.H2O [Cd(L)(H2O)2]6H2O [Zn(L)(H2O)2]H2O

1.49 1.39 1.34 1.33 1.33

1.29 1.34 1.35 1.33 1.40

1.21 1.22 1.22 1.215 1.21

1.20 1.21 1.214 1.217 1.22

0.95 — — — 0.99

0.95 — — — ___

— 1.91 1.91 1.82 2.00

— 1.91 2.00 1.82 1.88

— 1.92 1.93 1.82 1.87

— 1.92 1.94 1.82 1.87

— 1.98 2.05 — 1.99

1.96 MANO3

1.32

1.32

1.21

1.22

—

—

1.88

1.91

1.86

1.89

1.90

1.90 MANO3

1.33 1.32

1.33 1.32

1.22 1.22

1.23 1.23

— —

— —

2.2 2.1

2.2 2.1

2.26 2.05

2.26 2.05

4.45 2.4

Fig. 5. Molecular modeling of [Cu(L)(H2O)] (1).

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3.8. Antimicrobial activity The Schiff base, H2L, ligand and its metal complexes were evaluated for sensitive organisms which are: two strain Staphylococcus aureus and Bacillus subtilis as Gram-positive bacteria, Escherichia coli and Salmonella typhimurium as Gram-negative bacteria and Yeast: Candida albicans and Fungus: Aspergillus fumigatus. The activity of the Schiff base, H2L, ligand and its complexes against the used organisms are depicted in Table 6. The Schiff base, H2L, ligand showed no activity towards bacteria and fungus. From the results depicted in Table 6, we can conclude that: 1. Zn(II) complex 9 showed high activity towards Bacillus subtilis as Gram-positive bacteria and Candida albicans and Aspergillus fumigatus as fungus strain, while showed moderate activity towards Salmonella typhimurium and Escherichia coli as Gramnegative bacteria. 2. Fe(III) complex 6 showed high activity towards Staphylococcus aureus as Gram-negative bacteria and Escherichia coli as Gramnegative bacteria. 3. Cu(II) 2, Co(II) 4 and Cd(II) 8 complexes showed moderate activity towards Bacillus subtilis as Gram-positive bacteria, while Co(II) 4 and UO2(VI) 10 complexes showed moderate activity towards Salmonella typhimurium and Ni(II) complex 3 showed moderate activity towards Escherichia coli. 4. Cu(II) 1, Cu(II) 2, Ce(III) 7, and Fe(III) 6 complexes showed moderate activity towards Candida albicans as fungus strain. The increased activity upon chelation is attributed to the positive charge of metal partially shared with donor atoms present on ligand and possible p-electron delocalization over the whole chelate ring [57]. This, in turn, increases the lipophilic character

of the metal chelate and favors its permeation through the lipid layers of the bacterial membranes resulting in interference with normal process. Inhibition was found to increase with increasing concentration of metal complex [58]. 3.9. Molecular orbital calculations We were trying to assess the observed data at molecular level with the help of molecular modeling. This modeling program was commonly known as hyperchem 7.52. Molecular modeling had been successfully used to detect three dimensional arrangements of atoms in complexes by means of a semi-empirical molecular orbital calculations at the PM3 level provides by the hyperchem 7.52 program. Their utilization in the demonstration of molecular structure and structural properties data of the studied ligand and its complexes are presented in Tables 7 and 8, except VO(IV), Ce(III), and UO2(VI) complexes (5, 7, 10), respectively. The optimized structures of the metal complexes are shown in Figs. 5–8. The calculated energies of the Frontier orbitals, lowest unoccupied (ELUMO) and highest occupied molecular orbitals (EHOMO) and Egap (Egap = ELUMO – EHOMO) of the metal complexes are correlated with the current experimental data (Table 2) ELUMO/ eV = 32.638 + 0.0292 DCAO phenolic/cm1, R2 = 0.97, (n = 4 points except 3, 6, 9 complexes). The positive slope of the linear relationship reveals increasing of ELUMO (less stable complex) was accompanied by the blue shift of the stretching frequency of the CAO phenolic group. This emphasized by increasing the heat of formation (DHf) (more positive, less stable) with the increasing of the stretching frequency of the CAO phenolic as indicating from the positive slope of relationship of Hf versus DCAO phenolic, DHf/ kcal = 13468 + 12.21 DCAO phenolic/cm1, R2 = 0.95, (n = 5 points except 3 and 6 complexes). Consequently, Egap/eV = 5.54 + 0.052

Fig. 6. Molecular modeling of [Ni(L)] H2O (3).

O.M.I. Adly et al. / Journal of Molecular Structure 1054–1055 (2013) 239–250

249

Fig. 7. Molecular modeling of [Co(HL)(H2O)(NO3)]2.5H2O (4).

Fig. 8. Molecular modeling of [Fe(L)(H2O)(NO3)] 0.5CH3OHH2O (6).

DmC@O/cm1, R2 = 0.95 (n = 5 except 8 and 9 complexes). The positive slope of this linear relationship reveals increasing of Egap (high stable complex) accompanied by more splitting of the stretching frequency of the mC@O. In contrary, increasing of EHOMO was accompanied by less splitting of mC@O separation as indicated from the negative slope of the linear relationship between EHOMO versus DmC@O, EHOMO = 6.24–0.076 DmC@O, R2 = 0.97 (n = 5 points, except 2 and 8); The positive slope in the former relationships refer to the strong interaction of M(II) ion with the phenolic oxygen of the ligand which in turn increase of Egap values i.e., the separation between EHOMO and ELUMO increases (high stable). This emphasized

by the negative slope of the later relationship. Furthermore, the positive slope of the linear correlation of Egap versus the wavelength of the d–d transition in DMF solution, Egap/nm = 444.19 + 1.08 kmax/ nm, R2 = 0.99 (n = 3 except 3 and 4 complexes), this indicates red shift of the d–d transition with the increasing of the wavelength of the Egap energy. The interpretation given above further emphasized by the negative slope of the linear relationship of the calculated bond length of MAO(L) versus DmC@O; bond length MAO/Ao = 2.504 – 0.015 DmC@O /cm1, R2 = 0.97, (n = 4 points, except 1, 2, 9 complexes). This means that as the DmC@O separation frequency decreases the

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bond length between the phenolic oxygen and metal ion increases (elongated, weak bond strength), which agree well with the Gutmann’s bonds variation rule [29]. Moreover, the positive slope of relationship of DmC@O versus mMAO(L); mMAO(L) = 445.13 + 0.665 DmC@O/cm1, R2 = 0.91, (n = 4 except 3, 4, 9 complexes) further emphasized this interpretation revealing a dynamic redistribution of the electronic density on the whole system. 4. Conclusion The Schiff base (H2L) of 6-hydroxy-5-[N-(2-{[(1E)-1-(6-hydroxy-2,4-dioxo-3,4-dihydro-2H-1,3-thiazin-5yl)ethylidene]amino} ethyl)ethanimidoyl]-2H-1,3-thiazine-2,4(3H) dione, and its complexes with Cu(II), Ni(II), Co(II), VO(IV), Zn(II), Cd(II), Ce(III), Fe(III) and UO2(VI) have been synthesized and characterized. The ligand behaves as monobasic or dibasic tetradentate with N2O2 sites forming mononuclear complexes with different geometries. Cu(II) complex 1 form square-pyramidal geometry, Ni(II) form square planer geometry, Cu(II) 2, Co(II) 4, VO(IV) 5, Fe(III) 6, Cd(II) 8 and Zn(II) 9 complexes form octahedral geometry, while Ce(III) 7 and UO2(VI) 10 complexes in which the metal ions are octa-coordinated. Molecular parameters of the ligand and its metal complexes have been calculated and correlated with their experimental data. The antimicrobial activities of ligand and its metal complexes were screened and showed variable activities. Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.molstruc.2013. 09.037. References [1] T.T. Tidwell, Angew. Chem. Int. Ed. 47 (2008) 1016. [2] A. Shokrollahi, M. Ghaedi, M. Montazerozohori, A.H. Kianfar, H. Ghaedi, N. Khanjari, S. Noshadi, S. Joybar, E-J. Chemistry 8 (2011) 495. [3] A.H. Kianfar, L. Keramat, M. Dostani, M. Shamsipur, M. Roushani, F. Nikpour, Spectrochim. Acta Part A 77 (2010) 424. [4] A.H. Kianfara, S. Zargari, H.R. Khavasi, J. Iran. Chem. Soc. 7 (2010) 908. [5] R. Yuan, Y. Chai, D. Liu, D. Gao, J. Li, R. Yu, Anal. Chem. 65 (1993) 2572. [6] Y. Ohashi, Y. Bull, Chem. Soc. Jpn. 70 (1997) 1319. [7] S. Yamada, Coord. Chem. Rev. 1 (1966) 415. [8] A. Costes, F. Dahan, B. Donnadieu, M.-I. Fernandez-Garcia, M.-J. RodriguezDouton, Dalton Tarns. 19 (2003) 3776. [9] S. Kumar, D.N. Dahr, J Sci. Ind. Res. 68 (2009) 181. [10] P.J. McCarthy, R.J. Hovey, K. Ueno, A.E. Martell, J. Am. Chem. Soc. 77 (1955) 5820. [11] H.L. Singh, S. Varshney, A.K. Varshney, Appl. Organometal. Chem. 14 (2000) 212, and references therein. [12] M. Wang, L.F. Wang, Y.-Z. Li, Q.-X. Li, Z.-D. Xu, D.-M. Qu, Trans. Met. Chem. 26 (2001) 307. [13] N.N. Gulerman, S. Rollas, H. Erdeniz, M. Kiraj, J. Pharm. Sci. 26 (2001) 1.

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