Syntheses, characterization, biological activity and ...

Spectrochimica Acta Part A 81 (2011) 317–323

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Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy journal homepage: www.elsevier.com/locate/saa

Syntheses, characterization, biological activity and fluorescence properties of bis-(salicylaldehyde)-1,3-propylenediimine Schiff base ligand and its lanthanide complexes Ziyad A. Taha a,∗ , Abdulaziz M. Ajlouni a , Khader A. Al-Hassan b , Ahmed K. Hijazi c , Ari B. Faiq a a

Department of Applied Chemical Sciences, Jordan University of Science and Technology, Irbid 22110, Jordan Department of Chemistry, Faculty of Science, Yarmouk University, Irbid 1163, Jordan c Department of Chemistry, Faculty of Science, Jerash University, Jerash, Jordan b

a r t i c l e

i n f o

Article history: Received 16 March 2011 Received in revised form 12 May 2011 Accepted 13 June 2011 Keywords: Schiff base Lanthanides Luminescence Antibacterial activity

a b s t r a c t Eight new lanthanide metal complexes [LnL(NO3 )2 ]NO3 {Ln(III) = Nd, Dy, Sm, Pr, Gd, Tb, La and Er, L = bis-(salicyladehyde)-1,3-propylenediimine Schiff base ligand} were prepared. These complexes were characterized by elemental analysis, thermogravimetric analysis (TGA), molar conductivity measurements and spectral studies (1 H NMR, FT-IR, UV–vis, and luminescence). The Schiff base ligand coordinates to Ln(III) ion in a tetra-dentate manner through the phenolic oxygen and azomethine nitrogen atoms. The coordination number of eight is achieved by involving two bi-dentate nitrate groups in the coordination sphere. Sm, Tb and Dy complexes exhibit the characteristic luminescence emissions of the central metal ions attributed to efficient energy transfer from the ligand to the metal center. Most of the complexes exhibit antibacterial activity against a number of pathogenic bacteria. © 2011 Elsevier B.V. All rights reserved.

1. Introduction Schiff bases are considered as a very important class of organic compounds because of their ability to form stable complexes with many different transition metal and rare-earth metal ions in various oxidation states via N and O atoms [1–15]. They have the potential to be used in different areas such as electrochemistry, bioinorganic, catalysis, metallic deactivators, separation processes, and environmental chemistry. Moreover they are becoming increasingly important in the pharmacological, dye, and plastic industries as well as in the field of liquid crystal technology [16,17]. Since the first report of their metal complexes, simple di-, ter-, tetra-, and pentadentate Schiff base ligands have been extensively studied and used for the metal complexation. The tetradentate salen-type Schiff base ligands derived from salicylaldehyde and diamine complexes with transition metal ions have been synthesized and investigated using different chemical techniques. Their chemical analysis showed the formation of 1:1 [M:L] ratio complexes where the Schiff base ligands were coordinated to the metal ion in a tetradentate manner with N2 O2 donor sites of the two phenol-O and two azomethine-N [5–7]. In the last decades, the luminescent properties of lanthanides have attracted much attention for their wide applications in light

∗ Corresponding author. Tel.: +962 27201000x23640; fax: +962 27201071. E-mail address: [email protected] (Z.A. Taha). 1386-1425/$ – see front matter © 2011 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2011.06.018

emitting diodes (OLEDs), liquid crystal, fluoroimmunoassays, biophysics, laser technology, and optical telecommunication systems [18,19]. Mainly due to their very narrow emission bands, long excited-state lifetimes and large Stokes shifts. Since f–f electronic transitions of lanthanides in their “+3” oxidation state are Laporte forbidden, the direct photo-excitation of lanthanide ions is difficult [20]. Therefore, it is necessary to sensitize the lanthanide ions with chelating organic chromophores such as aromatic carboxylic acids, aromatic phenols, cryptand and heterocyclic ligands. Subsequently the absorbed light by the ligands can be transferred to the lanthanide ion (antenna effect) by an intra-molecular energy transfer [20]. The luminescent intensities of the lanthanide metal complexes are strongly dependent on the efficiency of the organic ligand to absorb UV light, the efficiency of energy transfer from the ligand to metal, and the efficiency of lanthanide metal luminescence. This work focuses on the use of a tetradentate salene-type ligand which has selective ability to coordinate to lanthanide ions thus protect them from deactivation caused by interaction with solvent molecules and enhance their luminescence by providing proper conjugate absorption groups suitable for energy transfer. Structurally characterized salen-type lanthanide complexes are rare. Depending on the preparative procedures, different compositions salen-type Ln(III) have been reported such as [Ln(H2 salen)(NO3 )3 ], [Ln2 (H2 salen)3 (NO3 )4 ](NO3 )2 (H2 O)3 , and [Ln2 (H2 salen)1.5 (NO3 )3 ]n [Ln2 (salen)3 ] [9–15]. In this context, we have studied the interaction of [Ln(NO3 )3 ·xH2 O] {Ln(III) = Nd, Dy, Sm, Pr, Gd, Tb, La, Er, x = 6

318

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except for Tb = 5} with bis-(salicyladehyde)-1,3-propylenediimine tetradentate Schiff base ligand L. The coordination behaviors have been investigated by correlating with their elemental analysis, thermal properties, 1 H NMR, FT-IR, UV–vis, and molar conductivity measurements. Moreover, the antimicrobial efficiency of these complexes has been screened against three different microorganisms. The photoluminescence properties of these complexes have been explored in solutions at room temperature through detailed photophysical investigation. 2. Experimental 2.1. Materials and methods [Ln(NO3 )3 ·6H2 O] {Ln = Sm, Nd, Er, Gd, Pr}, [Tb(NO3 )3 ·5H2 O] and [Dy(NO3 )3 ·xH2 O] were purchased from Sigma Aldrich Chemical company. Salicylaldehyde and 1,3-propyldiamine were purchased from Merck Schuchardt. All other solvents and reagents were of analytical grade purchased from TEDIA Company, OH, USA and used without further purification. The elemental analysis (C, H and N) was performed on Euro EA elemental analyzer 3000. The metal content of the complexes was determined by titration with EDTA [21]. Infrared spectra (IR) were recorded on a JASCO FT-IR model 470 spectrophotometer in the region of 4000–400 cm−1 using KBr pellets. The spectra were recorded at room temperature with 2 cm−1 resolution. The 1 H NMR spectra were recorded on a Bruker Avance 400 MHz spectrometer. The chemical shifts were measured in ppm in D6 -DMSO with tetramethylsilane as an internal standard. Molar conductivities were measured in dimethylformamide (DMF) solution, concentration of 10−4 M, at 25 ◦ C using WTW LF 318 conductivity meter equipped with WTW Tetracon 325 conductivity cell. Thermal analysis was performed on a PCT-2A thermobalance analyzer operating at a heating rate of 10 ◦ C/min in the range of 25 ◦ C up to 900 ◦ C under N2 . UV–vis spectra were recorded in ethanol solution, concentration of 10−5 M, at 25 ◦ C using a UV2401PC UV-visible spectrophotometer (Schimadzu Corporation). Fluorescence spectra (scanned from 200 to 700 nm, with a spectral resolution of 2.0 nm, slit widths ∼2.5 nm) were recorded on Edinburgh instrument model FS900SDT spectrometer with 1 cm quartz cell at constant room temperature. The light source and detectors were 450 W xenon lamp and R955 photomultiplier tube, respectively. 2.2. Synthesis of the Schiff base ligand L The Schiff base ligand L was prepared according to literature method [22]. Briefly, 2 equiv. of the salicylaldehyde were condensed with 1 equiv. of 1,3-propyldiamine in refluxing ethanol. Upon cooling, a crude yellow product was obtained. Then, the yellow product was collected by filtration, washed with cold ethanol,

and air-dried. Finally, the product was recrystallized from hot ethanol. The melting point, 1 H NMR and 13 C NMR were consisting with literature reports. 2.3. Synthesis of the Ln(III) complexes A 1.0 mmol (282.3 mg) of the ligand L was dissolved in 10 mL chloroform. To this solution, a solution of 1.0 mmol (435.01 mg) Pr(NO3 )3 ·6H2 O in 10 mL ethyl acetate was drop wise added. The reaction mixture was stirred for 2 h at room temperature. Then, the yellow precipitate was filtered, washed several times with ethyl acetate and chloroform, and after that dried for 24 h under vacuum at room temperature. All other Ln(III) complexes were prepared in a similar manner. 2.4. Biological activity The isolated bacteria used in this study were obtained from the Central Laboratories, Jordan Ministry of Health. The clinical isolates (Shigella dysenteriae, Proteus vulgaris, and Pseudomonas aeruginosa) were grown in Muller Hinton Broth media for 24 h at 37 ◦ C. The biological activity of the ligand L and its Ln(III) complexes was estimated by a Minimum Inhibitory Concentration (MIC, ␮g/mL) using a microbroth dilution method and following the guidelines of Hannan [23]. A stock solutions of the ligand L and its Ln(III) complexes in DMSO were prepared according to NCCLS guidelines [24]. The in vitro MIC was carried out in standard sterile 96 well flat bottom micro-titer plates. The layout was designed so that each row covered the final antimicrobial dilution of 500–0.5 ␮g/mL with one control well. To each well, a 40 ␮L of the selected complex at the correct concentration was added and the control well was loaded with a 40 ␮L of DMSO solvent. Then a 150 ␮L of Muller Hinton media was added to all wells, followed by a 10 ␮L of the bacterial culture at approximately 108 cfu/cm3 , giving a final concentration of the bacteria of approximately 5 × 107 cfu/cm3 . The plates were sealed and incubated at 37 ◦ C under atmospheric conditions. After 24 hr incubation the micro-titer plates were read using a ELIZA UV-visible spectrometer. The minimal concentration that had optical density less than the control was defined as the MIC. 3. Results and discussions In the present investigations, the ligand L and its new eight Ln(III) complexes were synthesized and characterized by different physical and chemical techniques. Scheme 1 summarizes the multi-step procedure leading to the target complexes. The ligand L was synthesized by a conventional one-step condensation of salicylicaldehyde and 1,3-propyldiamine and characterized by 1 H NMR (Fig. 1SM), IR, 13 C NMR (Fig. 2SM), TGA (Fig. 3SM) and elemental analysis [22].

O 2

N H

OH

+ H2N

NH2

Ethanol Ref lux 2h

N

OH

HO L

Ln(NO3)3.xH2O ethylacetate R. T chloroform [LnL(NO3)2]NO3 Scheme 1. Synthetic route for Schiff base ligand L and its Ln(III) complexes.

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Molar conductivity (m ) values for all complexes in DMF solutions at 25 ◦ C, tabulated in Table 1, are in the range of 109–125 −1 cm2 mol−1 reported for 1:1 electrolytes [25]. These values suggest that two nitrate groups are coordinated to the Ln(III) ions. 3.3. Thermal gravimetric analysis (TGA) The water content was determined by the thermogravimetric analysis (TGA). The TGA results showed that the ligand L is thermally stable in the temperature range 25–225 ◦ C and its decomposition starts at 225 ◦ C and finishes at 327 ◦ C with one decomposition step, Fig. 3SM. All Ln(III) complexes are thermally stable up to 200 ◦ C indicating the absence of lattice or coordinated water and solvent molecules, Fig. 3SM. The thermal decomposition of these complexes proceeds with four decomposition steps. The first and second steps correspond to the loss of two and one nitrate species from the inner and the outer coordination sphere of these complexes, respectively. Third and forth steps involve the loss of the ligand L and the formation of Ln2 O3 as the final residue

3.5.

1H

NMR characterization

To assess the binding nature of the ligand L to the Ln(III) ions, 1 H NMR spectra of the ligand L and its La(III) complex were scanned in the range 0–16 ı ppm. The ligand L showed two singlet peaks at 13.51 and 8.57 ı ppm, Fig. 1SM, were assigned to the phenolic hydroxyl protons (Ar–OH) and the azomethine protons (–CH N–), respectively [22]. The quintet and the triplet peaks at 3.71 and 2.50 ı ppm were attributed to the methylene protons, –CH2 –C–CH2 – and –C–CH2 –C–, respectively. In addition, a set of multiplet was observed in the range of 6.88–7.31 ı ppm assigned to the protons of the aromatic rings ( C–H). In the 1 H NMR spectrum of the La(III) complex, the peaks corresponding to the phenolic hydroxyl and azomethine protons are

(i)

3.4. Infrared characterization

-1385

-1320

-1483

-1651

Absorbance

(a)

(b)

(c)

1800

1700

1600

1500

1400

1300

1200

1100

Wavenumber/cm-1 (ii)

(a)

-816

-579

-1028

Absorbance

The binding mode of the ligand L to Ln(III) ions in these complexes was studied by comparing the IR spectrum of the ligand L with that of its Ln(III) complexes. The IR spectrum of the ligand L shows two strong absorption bands at 1635 and 1280 cm−1 , which are attributed to the (C N) of the azomethines and (Ar–O) of the phenolic hydroxyl substituent, respectively, Fig. 1a [22]. The infrared spectra of all Ln(III) complexes displayed the ligand L characteristic bands with various shifts due to complex formation, Table 2. The IR spectral shifts of different complexes are similar, indicating similar structures of the Ln(III) complexes. The IR spectrum of the Tb(III) complex shows a significant increase in the C–O stretching frequency (10 cm−1 ) compared to the free ligand L, Fig. 1b. This suggest that the coordination to the Tb(III) ion occurs through the oxygen atoms of the hydroxyl benzene of the ligand L [14]. This was further confirmed by the appearance of a medium intensity band at 579 cm−1 which may be assigned to (Tb–O) vibration, Fig. 1b. In comparison with the free ligand L, the (C N) band in the Tb(III) complex shifted to 1651 cm−1 . This blue shift (16 cm−1 ) indicates a stronger double bond character of the iminic bonds and participation of the azomethines nitrogen atoms in the coordination [14,22], Fig. 1b. This data was further supported by the appearance of a medium intensity band at 411 cm−1 which could be assigned to (Tb–N) vibration, Fig. 1b. In addition, the IR spectrum of the free ligand L exhibits a broad band in the range (3500–3340 cm−1 ) centered at 3445 cm−1 due to the stretching vibration of phenolic groups (O–H) perturbed by intramolecular hydrogen bonding (O–H· · ·N) between the phenolic hydrogen

-411

3.2. Molar conductivity measurements

-1280

The results of the elemental analysis of the ligand L and its Ln(III) complexes, listed in Table 1, are in good agreement with those calculated based on molecular formula proposed.

-1290

3.1. Elemental analysis

and the azomethines nitrogen atoms [26,27]. In the spectra of the Tb(III) complexes, this band shifted to 3433 cm−1 indicating that the hydroxyl oxygen is coordinated to the Tb(III) ion without a proton displacement. The 1 H NMR spectrum of the Ln(III) complexes also shows the presence of the OH signal confirming the IR data. The most important aspect concerning the infrared spectra of all Ln(III) complexes is the possibility of characterizing the different coordination modes of the nitrate groups to the lanthanide metals which is usually detected by the differences in the two bands |4 − 1 |. The infrared spectrum of the Tb(III) complex displays several intense bands at 1483 cm−1 (1 ), 1029 cm−1 (2 ), 816 cm−1 (3 ) and 1317 cm−1 (4 ) assigned to the coordinated nitrate ions (C2v ), Fig. 1. The difference between 4 and 1 is approximately 166 cm−1 suggesting that the coordinated NO3 − ions in the complex act as a bidentate ligand [28,29]. The (o ) free nitrate appears at 1385 cm−1 in the spectrum of the complex agrees with the results of the conductivity experiment. Similar changes were observed for all other Ln(III) complexes upon the complexation with the ligand L.

-1635

All Ln(III) complexes were synthesized by treating [Ln(NO3 )3 ·xH2 O] with the ligand L which yielded a series of complexes correspond to the formula of [LnL(NO3 )2 ]NO3 . All complexes were stable in air, non-hygroscopic powders, soluble in DMSO and DMF, slightly soluble in methanol, ethanol, ethyl acetate, chloroform and benzene, and insoluble in water and diethyl ether.

319

(b)

(c)

1100

1000

900

800

700

600

500

400

-1

Wavenumber/cm

Fig. 1. Infrared spectra of (a) Schiff-base ligand L; (b) [TbL(NO3 )2 ]NO3 ; (c) [DyL(NO3 )2 ]NO3 in the regions: (i) 1800–1100 cm−1 ; (ii) 1100–400 cm−1 .

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Table 1 Analytical data and molar conductance values for the Schiff base ligand L and its Ln(III) complexes. Complex

C (%) found (calc.)

H (%) found (calc.)

N (%) found (calc.)

L [NdL(NO3 )2 ]NO3 [SmL(NO3 ) 2 ]NO3 [DyL(NO3 )2 ]NO3 [LaL(NO3 )2 ]NO3 [ErL(NO3 )2 ]NO3 [PrL(NO3 )2 ]NO3 [TbL(NO3 ) 2 ]NO3 [GdL(NO3 )2 ]NO3

72.20 (72.24) 32.42 (32.85) 31.91 (32.53) 31.36 (31.91) 32.85 (33.13) 32.15 (31.67) 33.34 (33.02) 32.27 (32.09) 32.22 (32.17)

6.45 (6.32) 3.12 (3.08) 2.93 (3.05) 3.14 (2.99) 3.05 (3.11) 2.96 (2.97) 3.17 (3.10) 3.15 (3.01) 3.12 (3.02)

10.03 (9.91) 11.53 (11.27) 11.32 (11.16) 11.36 (10.95) 11.69 (11.36) 11.31 (10.86) 11.12 (11.33) 11.42 (11.01) 11.36 (11.04)

Ln (%) found (calc.) – 23.60 (23.20) 30.00 (29.31) 24.77 (25.40) 21.85 (22.54) 25.27 (25.95) 22.00 (22.79) 25.22 (24.98) 28.59 (28.99)

Yield (%)

m (S cm2 mol−1 )

– 71 73 71 62 98 98 98 90

– 117.01 119.17 115.92 117.25 108.77 104.36 114.42 115.18

Table 2 Major infrared spectral data for the Schiff base ligand L and its Ln(III) complexes (cm−1 ). Complex

(O–H)

(C N)

(Ar–O)

L [NdL(NO3 )2 ]NO3 [SmL(NO3 )2 ]NO3 [DyL(NO3 )2 ]NO3 [LaL(NO3 )2 ]NO3 [ErL(NO3 )2 ]NO3 [PrL(NO3 )2 ]NO3 [TbL(NO3 )2 ]NO3 [GdL(NO3 )2 ]NO3

3445 3433 3435 3433 3433 3435 3436 3433 3434

1635 1650 1651 1650 1650 1651 1651 1651 1651

1280 1287 1289 1291 1289 1288 1287 1290 1288

(NO3 − ) 1

2

3

4

1 –4

o

– 1480 1481 1482 1481 1484 1481 1483 1483

– 1029 1028 1029 1030 1029 1029 1029 1028

– 817 816 816 818 815 817 816 816

– 1320 1316 1311 1320 1314 1319 1317 1319

– 160 165 171 161 170 162 166 164

– 1384 1384 1385 1385 1385 1385 1385 1385

slightly up-field shifted compared to the free ligand L. This shift indicates the involvement of nitrogen atoms of the azomethine and oxygen atoms of the hydroxyl benzene in the coordination without deprotonation which is consistence with the IR data discussed in the previous section. Moreover, there was no noticeable change in the peak positions corresponding to aromatic and methylenic protons.

(M–O)

(M–N)

– 577 578 580 575 580 577 579 578

– 410 410 411 410 412 413 411 413

groups which is overlapped with intramolecular charge transfer from the phenyl ring [7,22]. The similarity of UV–vis absorption spectra of all investigated complexes indicates similar structures of the Ln(III) complexes, Fig. 2. It is important to recall that lanthanides ion do not appreciably contribute to the spectra of their complexes since f-f transitions are Laporte-forbidden and very weak (extinction coefficients ε < 1 M−1 cm−1 ) [20]. On the other hand, charge transfer

3.6. Electronic spectroscopy 3.6.1. UV–vis spectroscopy The UV–vis absorption spectra values of the maximum absorption wavelength (max ) and the corresponding molar absorptivities (ε) for the ligand L and its Ln(III) complexes (10−5 M in ethanol) are listed in Table 3. Fig. 2a shows the absorption spectrum of the ligand L which is characterized by four main absorption bands in the region 200–500 nm. The first two high intensity bands observed at max = 215 nm and max = 255 nm are attributed to the ␲–␲* transitions of the benzene rings [19,30,31]. The third absorption band at max = 316 nm may correspond to the ␲–␲* transition of the azomethine group C N while the last band at max = 407 nm may correspond to the n–␲* transitions associated with the azomethine 0.15

Absorbance

b c d e

Table 3 The UV–vis absorption bands (max ) and molar absorptivities (ε) of the Schiff base ligand L and its Ln(III) complexes in ethanol solution at the room temperature. Compound

max (nm)

ε × 104 (cm2 mol−1 )

Band assignment

Ligand

215, 255, 316 407

3.55, 1.97, 0.677 0.125

␲ → ␲* n → ␲*

[ErL(NO3 )2 ]NO3

220, 255, 319 241, 355 410

4.24, 2.17, 0.686 2.26, 0.374 0.1991

␲ → ␲* LMCT n → ␲*

[GdL(NO3 )2 ]NO3

218, 255, 320 241, 355 409

3.88, 1.92, 0.639 2.06, 0.321 0.1780

␲ → ␲* LMCT n → ␲*

[NdL(NO3 )2 ]NO3

218, 255, 319 241, 356 409

5.02, 2.26, 0.65 2.51, 0.39 0.0782

␲ → ␲* LMCT n → ␲*

[PrL(NO3 )2 ]NO3

219, 255, 320 242, 355 409

2.99, 1.9, 0.67 2.0, 0.281 0.1467

␲ → ␲* LMCT n → ␲*

[SmL(NO3 )2 ]NO3

219, 256, 320 240, 356 410

3.28, 6.09, 0.527 1.95, 0.282 0.1167

␲ → ␲* LMCT n → ␲*

[TbL(NO3 )2 ]NO3

221, 255, 320 240, 356 409

3.86, 1.81, 0.577 1.95, 0.296 0.1587

␲ → ␲* LMCT n → ␲*

[LaL(NO3 )2 ]NO3

222, 255, 320 239, 355 410

3.78, 1.70, 0.551 2.61, 0.478 0.1176

␲ → ␲* LMCT n → ␲*

[DyL(NO3 )2 ]NO3

223, 254, 320 241, 355 410

5.36, 3.07, 0.702 3.09, 0.393 0.1000

␲ → ␲* LMCT n → ␲*

0 350

375

400

425

450

a 200

250

300

350

400

450

500

Wavelength/ nm Fig. 2. UV–vis absorption spectra of (a) Schiff base ligand L; (b) [TbL(NO3 )2 ]NO3 ; (c) [PrL(NO3 )2 ]NO3 ; (d) [DyL(NO3 )2 ]NO3 ; (e) [ErL(NO3 )2 ]NO3 in ethanol solution at the room temperature.

Z.A. Taha et al. / Spectrochimica Acta Part A 81 (2011) 317–323

Fluorescence Intensity

b

320

c d e f

360

400

440

480

520

560

Wavelength/ nm Fig. 3. Emission spectra of (a) Schiff base ligand L; (b) [GdL(NO3 )2 ]NO3 ; (c) [LaL(NO3 )2 ]NO3 ; (d) [ErL(NO3 )2 ]NO3 ; (e) [PrL(NO3 )2 ]NO3 ; (f) [NdL(NO3 )2 ]NO3 in DMF solution (1.0 × 10−5 M) at the room temperature. Emission spectra are obtained with exc = 250 nm.

bands involving lanthanide orbitals are typically neither observed in the near-UV nor visible spectral regions. Hence the absorption bands of lanthanide complexes are completely attributable to ligand-centered transitions, although some perturbation is observable upon complexation. The absorption bands located at around 215 nm, 316 nm, and 407 nm are slightly shifted to lower absorption frequencies on complexation. The complexation was further supported by the appearance of two new absorption bands at 241 nm and 355 nm, which were absent for the free ligand L and Ln(NO3 )3 . The absorption band observed at 241 was attributed to the electron transfer between lanthanide metal ions and the coordinated ligand [15]. While the one at 355 may be assigned to ␲–␲* electronic transitions of C N groups coupled with charge transfer from the ligand to metal ion [32]. The band shifts, the appearance of the new two bands and the change in ε indicate the formation of Ln(III) complexes.

exhibits broad fluorescence band centered at 433 nm attributed to ␲ → ␲* transitions, Fig. 3a [33,34]. Inspection of emission spectra (exc = 250 nm) shown in Fig. 4 together with the data listed in Table 4, indicate that [SmL(NO3 )2 ]NO3 , [DyL(NO3 )2 ]NO3 and [TbL(NO3 )2 ]NO3 complexes exhibit the characteristic emission spectra of the Sm(III), Tb(III) and Dy(III) ions, respectively. This indicates that the ligand L is a good chelating organic chromophore and can be used to absorb and transfer energy to the Ln(III) ions. The sensitized emission spectrum of the [SmL(NO3 )2 ]NO3 complex displays three luminescence bands at 17,637, 16,667, and 15,456 cm−1 corresponding to the 4 G5/2 → 6 H5/2 at 567 nm, 4G 6 4G 6 5/2 → H7/2 at 600 nm, and 5/2 → H9/2 at 647 nm transitions, respectively, Fig. 4a. The intensity sequence of the peaks is I4G5/2 → 6H9/2 > I4G5/2 → 6H7/2 > I4G5/2 → 6H5/2 similar to that previously reported for samarium bis-diketonate complexes [35]. For the Tb(III) complex, the emission spectrum excited at 250 nm shows four luminescence bands at 20,408, 18,282, 17,007, and 16,129 cm−1 corresponding to the 5 D4 → 7 F6 at 490 nm, 5 D4 → 7 F5 at 547 nm, 5 D4 → 7 F4 at 588 nm, and 5 D4 → 7 F3 at 620 nm transitions, respectively, Fig. 4b (dotted line). Among these transitions, the 5 D4 → 7 F5 transition exhibits the strongest green emission and

4

(a) Fluorescence Intensity

a

321

4

G

4

400

450

G

500

6 5/2

H

G

6

H

5/2

9/2

6

H

5/2

7/2

5/2

550

600

650

700

Wavelength/nm

Table 4 Luminescence spectra data of the Schiff base ligand L and its Ln(III) complexes in ethanol at the room temperature. ex (nm)

Ligand L [LaL(NO3 )2 ]NO3 [GdL(NO3 )2 ]NO3 [NdL(NO3 )2 ]NO3 [PrL(NO3 )2 ]NO3 [ErL(NO3 )2 ]NO3

250 250 250 250 250 250

[SmL(NO3 )2 ]NO3

[TbL(NO3 )2 ]NO3

[DyL(NO3 )2 ]NO3 a b

Broad. Sharp.

250

250

250

 (cm−1 )

em (nm)

 (cm−1 )

7

F

4

5

F

4

6

5

D

7

F

4

4

5

D

7 4

F

3

23,095 21,930 21,459 22,989 22,779 22,831

␲ → ␲* ␲ → ␲* ␲ → ␲* ␲ → ␲* ␲ → ␲* ␲ → ␲*

a

40,000

433 567b 600b 647b

23,095 17,637 16,667 15,456

␲ → ␲* 4 G5/2 → 6 H5/2 4 G5/2 → 6 H7/2 4 G5/2 → 6 H9/2

40,000

433a 490b 547b 588b 620b

23,095 20,408 18,282 17,007 16,129

␲ → ␲* 5 D4 → 7 F6 5 D4 → 7 F5 5 D4 → 7 F4 5 D4 → 7 F3

40,000

433a 479b 574b

23,095 20,876 17,422

␲ → ␲* 4 F9/2 → 6 H15/2 4 F9/2 → 6 H13/2

i ii

Assignment

433a 456a 466a 435a 439a 438a

40,000 40,000 40,000 40,000 40,000 40,000

D

D

7

5

400

450

500

550

600

650

700

650

700

Wavelength/nm

(c) Fluorescence Intensity

Compound

5

(b) Florescence Intensity

3.6.2. Luminescence spectroscopy Fig. 3 shows the emission spectra of the ligand L and its Ln(III) complexes (1.0 × 10−5 M in ethanol solution) were recorded at the room temperature. The details of the emission characteristic of the ligand L and its Ln(III) complexes are listed in Table 4. The emission spectrum of the free ligand L

400

4

450

F

4

6 9/2

H

F

15/2

500

6 9/2

550

H

13/2

600

Wavelength/nm Fig. 4. Emission spectra of the complexes: (a) [SmL(NO3 )2 ]NO3 , exc = 250 nm; (b) [TbL(NO3 )2 ]NO3 , (i) exc = 315 nm, (ii) exc = 250 nm; (c) [DyL(NO3 )2 ]NO3 , exc = 250 nm, in ethanol solution (1.0 × 10−5 M) at the room temperature.

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25000

647 nm

600 nm

567 nm

574 nm

479 nm

Ligand Fluorescence

Absorption

547 nm

490 nm

10000

588 nm

15000 620 nm

Energy/ cm

-1

20000

5000

0 Ground state of Tb(III)

Ground state of the ligand

Ground state of Dy(III)

Ground state of Sm(III)

Scheme 2. The probable mechanism of intramolecular energy transfer between the Schiff base ligand L and the Tb(III), Sm(III), and Dy(III) ions.

5D 4

→ 7 F6 transition shows the second strongest blue emission. The emission spectrum of Tb(III) excited at 315 nm shows similar luminescence bands as the one excited at 250 nm which indicates that the two excitation bands are the effective energy sensitizers for the luminescence of Ln(III) ions, Fig. 4b. For Dy(III) complex, the emission spectrum shows two luminescence bands at 20,876 and 17,422 cm−1 corresponding to the blue emission 4 F9/2 → 6 H15/2 at 479 nm and 4 F9/2 → 6 H13/2 at 574 nm transitions, respectively, Fig. 4c. The weak broad emission band in the spectra of Sm(III), Tb(III) and Dy(III) complexes is attributed to the fluorescence of the ligand L, however this band is almost invisible for the case of Sm(III) complex. This indicates that the ligand L to metal energy transfer process in the Sm(III) complex is more efficient than that in Tb(III) or Dy(III) complexes. This may be due to the better matches between the lowest triplet state energy level (TL ) of the ligand L and the lowest resonance energy level of the Sm(III) (4 G5/2 , 17,900 cm−1 ) ion as compared to that of Tb(III) (5 D4 , 205,000 cm−1 ) and Dy(III) (4 F9/2 , 21,000 cm−1 ) ions. It must be here mentioned that the small energy gap could result in a back-energy transfer process from the excited resonance levels of the Tb(III) and Dy(III) to the TL of the ligand L, therefore leading a depression in the luminescence output of the Tb(III) and Dy(III) ions [36–38]. Although different paths have been suggested for the energy transfer from TL of the ligand L to the resonance state of Ln(III) ions in lanthanide complexes. The favorite mechanism involves strong absorption of UV energy that excites the ligand L to the excited singlet state (SL ), followed by an energy migration via nonradiative intersystem crossing to the TL . The energy is then transferred intramolecularly from the TL to a resonance state of the Ln(III) ion, from which the luminescence in the visible region occurs. To luminescence, the TL level must be nearly equal or lie above the resonance energy level of the Ln(III) ions. If the energy gap between the TL and the lowest resonance energy level of the Ln(III) ion is low then the back transfer from the lanthanide ion to the ligand L can occur which reduces the intensity of the luminescence. Our results of the luminescence experiments were consistent with the proposed mechanism for intramolecular energy transfer shown in Scheme 2. For the case of [LaL(NO3 )2 ]NO3 , La(III) has no 4f electron and no excited states below the TL , while [GdL(NO3 )2 ]NO3 , Gd(III) posses a relatively stable 4f shell and the lowest-lying excited state 6 P7/2 located at about 32,000 cm−1 , which is much higher in energy than

the SL and TL of the ligand L [35]. Therefore, the energy absorbed by the ligand L cannot be transferred to either La(III) or Gd(III) ions by an intramolecular energy transfer process, however relaxes through its own lower energy levels. Therefore the emission bands observed in the emission spectra of La and Gd complexes at 456 nm and 466 nm, respectively, are due to the fluorescence of the ligand L, Fig. 3b and c. The red shift in the fluorescence band of the ligand in Ln(III) complexes compared with the free ligand L can be attributed to coordination of Ln(III) ions to the ligand L, which resulting in increasing of the delocalization of electrons and reducing the energy gaps between the ␲ → ␲* molecular orbitals of the ligand L [34,39]. Furthermore, the emission spectra of Nd, Pr, and Er complexes exhibit a single broad emission band at 435 nm, 439 nm, and 438 nm, respectively, which are attributed to ␲ → ␲* transitions of the ligand L, Fig. 3. The appearance of the strong fluorescence characteristic of the ligand L and the absence of the characteristic emission bands of Nd(III), Er(III) and Pr(III) ions is probably due to the large energy gap between the TL of the ligand and there lowest resonance energy levels. 3.7. Antibacterial activity The results of the bactericidal screening of the ligand L and its Ln(III) complexes are recorded in Table 5. The obtained results reflect that the ligand L has a good activity against Pseudomonas aereuguinosa and S. dysenteriae and does not show any activity against P. vulgaris. Most of the tested Ln(III) complexes possessed a high antibacterial activity against the Gram-negative bacteria tested. It was observed from the results that the activity of the Ln(III) complexes in most cases is higher than that of the corresponding ligand L. Enhancement in the activity upon complexation can be explained on the basis of chelation theory [40–42]. According to the chelation theory, the polarity of the Ln(III) ion is reduced by coordinating with ligand, which in turn increases the lipophilic character of the central metal atom. The lipophilic nature of Ln(III) enhanced its penetration into the lipid membrane of the microorganism cell wall, therefore raising the activity of the complex and more restriction on further growth of the organism. Furthermore, it has been suggested that some functional groups such as azomethine or hydroxyl groups present in these complexes play an important role in antibacterial activity [42,43]. Comparing the biological activity of the ligand L and its Ln(III) complexes with the standard Cephalexin

Z.A. Taha et al. / Spectrochimica Acta Part A 81 (2011) 317–323 Table 5 Minimum Inhibitory Concentration (MIC) of the Schiff base ligand L and its Ln(III) complexes against test bacteria. Compound

Ligand L [NdL(NO3 )2 ]NO3 [SmL(NO3 )2 ]NO3 [ErL(NO3 )2 ]NO3 [LaL(NO3 )2 ]NO3 [TbL(NO3 )2 ]NO3 [GdL(NO3 )2 ]NO3 [PrL(NO3 )2 ]NO3 [TbL(NO3 )2 ]NO3 Cephalexin Cephradine

MIC (␮g/mL) Shigella dysenteriae

Pseudomonas aeruguinosa

Proteus vulgaris

8 N 4 N 4 4 4 4 4 8 16

4 N 4 32 32 4 4 4 2 125 125

N 8 N 4 4 N 8 8 4 125 125

N = none active compound.

and Cephradine (antibacterial agent) reveals that most of the tested compounds are more potent than these standard antibiotics. 4. Conclusions In this work the tetradentate Schiff base ligand L and its corresponding lanthanide complexes [LnL(NO3 )2 ]NO3 are synthesized and characterized. It is concluded from analytical and spectral data that the ligand L is a tetradentate chelate and coordinates to the central Ln(III) ion by the two imine nitrogen atoms and the two phenolic oxygen atoms with 1:1 stoichiometry. IR spectral, conductivity and thermal data showed that two nitrate groups are bound in a bidentate manner to the central Ln(III) ions in the coordination sphere while another nitrate is in the outer coordination sphere in complexes. Hence coordination number eight is suggested for the metal ion in these lanthanide(III) nitrate complexes. The fluorescence date of the free ligand L and its Ln(III) complexes reveal that the Sm(III), Tb(III) and Dy(III) complexes exhibit characteristic luminescence of Sm(III), Tb(III) and Dy(III) ions, which indicates that the ligand L is a good organic chelator to absorb and transfer energy to Sm(III), Tb(III) and Dy(III) ions. According to the data discussed above it is obvious that the energy gap between the TL and the lowest excited state level of Sm(III) favor to the energy transfer process for Sm. Thus, Sm, Tb and Dy complexes can be a candidate as a luminescent in supramolecular nanodevice and laser materials. The antimicrobial activity results show that most of the synthesized complexes possess a good antibacterial activity against Gram-negative bacteria tested and the microbial activity of the complexes in most cases is higher than that of the corresponding ligand. Acknowledgment The authors are grateful to the Deanship of Scientific Research at Jordan University of Science and Technology for the financial support of this work, grant number (48/2008). Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.saa.2011.06.018.

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References [1] N.F. Choundhary, N.G. Connelly, P.B. Hitchcock, G.J. Leigh, J. Chem. Soc. Dalton Trans. 24 (1999) 4437–4446. [2] Z.L. You, L.L. Tang, H.L. Zhu, Acta Cryst. E 61 (2005) m36–m38. [3] L.B. Steven, D.R. John, H. Gloria, V. Meera, L. Charles, Inorg. Chem. 40 (2001) 972–976. [4] S. Chattopadhyay, M.G.B. Drew, A. Ghosh, Eur. J. Inorg. Chem. 10 (2008) 1693–1701. [5] S.A. Abdallah, G. Mohamed, M. Zayed, M. El-Ela, Spectrochim. Acta A 73 (2009) 833–840. [6] P.A. Vigato, S. Tamburini, Coord. Chem. Rev. 248 (2004) 1717–2128. [7] K. Karaoglu, T. Baran, K. Serbest, M. Er, I. De˘girmencio˘glu, J. Mol. Struct. 922 (2009) 39–45. [8] S.L. Barnholtz, J.D. Lydon, G. Huang, M. Venkatesh, C.L. Barnes, A.R. Ketring, S.S. Jurisson, Inorg. Chem. 40 (2001) 972–976. [9] Y.X. Ceng, M.Q. Jin, H.W. Shu, Progress in Coordination Chemistry, 251, vol. 8, Higher Education Press, Beijing, 2000, p. 17. [10] W. Radecka-Paryzek, I. Pospieszna-Markiewicz, M. Kubicki, Inorg. Chim. Acta 360 (2007) 488–496. [11] D. Wei, Y.I. Jun-Na, L.I. Wei-Sheng, W.U. Ya-Wen, Z. Jiang-Rong, W. Da-Qi, Inorg. Chem. 10 (2007) 105–108. [12] J. Chakraborty, S. Thakurta, G. Pilet, R.F. Ziessel, L.J. Charbonnière, S. Mitra, Eur. J. Inorg. Chem. 26 (2009) 3993–4000. [13] M.T. Kaczmarek, I. Pospieszna-Markiewicz, M. Kubicki, W. Radecka-Paryzek, Inorg. Chem. Commun. 7 (2004) 1247–1249. [14] W.B. Sun, P.F. Yan, G.M. Li, H. Xu, J.W. Zhang, J. Solid State Chem. 182 (2009) 381–388. [15] W. Xie, M.J. Heeg, P.G. Wang, Inorg. Chem. 38 (1999) 2541–2543. [16] S. Kumar, D.N. Dhar, P.N. Saxena, J. Sci. Ind. Res. 41 (1982) 501–506. [17] M.D. Deshmukh, A.G. Doshi, Orient. J. Chem. 11 (1995) 86–96. ˝ (Eds.), Lanthanide Probes in Life, Chem[18] J.F. Desreux, in: G.R. Choppin, J.C. Bunzli ical and Earth Sciences, Elsevier, Amsterdam, 1989. [19] L. Guo, S. Wu, F. Zeng, J. Zhao, Eur. Polym. J. 42 (2006) 1670–1675. [20] C. Görller-Walrand, K. Binnemans, Spectral intensities of f–f transitions, in: K.A. Gschneidner Jr., L. Eyring (Eds.), Handbook on the Physics Chemistry of Rare Earths, vol. 25, North-Holland Publishers, Amsterdam, 1998, pp. 101–264 (Chapter 167). [21] F.J. Welcher, The Analytical Use of EDTA 4, Van Nostrand, New York, 1965. [22] T.A. Youssef, J. Coord. Chem. 61 (2008) 816–822. [23] C.P. Hannan, International research programme on comparative mycoplasmology, Vet. Res. 31 (2000) 373–395. [24] G.L. Woods, J.A. Washington, Antibacterial susceptibility tests: dilution and disk diffusion methods, in: P.R. Murray, E.J. Baron, M.A. Pfaller, F.C. Tenover, R.H. Yolken (Eds.), Manual of Clinical Microbiology, 6th ed., American Society for Microbiology, Washington, DC, 1995, pp. 1327–1341. [25] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81–122. [26] M. Salavati-Niasari, Z. Salimi, M. Bazarganipour, F. Davar, Inorg. Chim. Acta 362 (2009) 3715–3724. [27] D.N. Kumar, B.S. Garg, Spectrochim. Acta Part A 64 (2006) 141–147. [28] W. Wang, Y. Huang, N. Tang, Spectrochim. Acta A 66 (2007) 1058–1062. [29] Y.L. Zhang, W.W. Qin, W.S. Liu, M.Y. Tan, N. Tang, Spectrochim. Acta A 58 (2002) 2153–2157. [30] R.C. Felicio, E.T.G. Canalheiro, E.R. Dockal, Polyhedron 20 (2001) 261–268. [31] S. Zolezzi, A. Decinti, E. Spodine, Polyhedron 18 (1999) 897–904. [32] M.M. Moawad, W.G. Hanna, J. Coord. Chem. 55 (2002) 439–457. [33] Y. Ho, W.Y. Yu, K.K. Cheung, C.M. Che, J. Chem. Soc. Dalton Trans. 10 (1999) 1581. [34] J.X. Li, Z.X. Du, J. Zhou, H.Q. An, S.R. Wang, B.L. Zhu, S.M. Zhang, S.H. Wu, W.P. Huang, Inorg. Chim. Acta 362 (2009) 4884–4890. [35] Y.X. Zheng, L.S. Fu, Y.H. Zhou, J.B. Yu, Y.N. Yu, S.B. Wang, H.J. Zhang, J. Mater. Chem. 12 (2002) 919–923. [36] W.N. Wu, N. Tang, L. Yan, Spectrochim. Acta A 71 (2008) 1461–1465. [37] M. Latva, H. Takalo, V.-M. Mukkala, C. Matachescu, C.J. Rodriguez-Ubis, J. Kankare, J. Lumin. 75 (1997) 149–169. [38] F. Gutierrez, C. Tedeschi, L. Maron, J.P. Daudey, R. Poteau, J. Azema, P. Tisnès, C. Picard, J. Chem. Soc. Dalton Trans. 9 (2004) 1334–1347. [39] S.J. Dunne, L.A. Summers, E.I.V. Nagy-Felsobuki, J. Heterocycl. Chem. 29 (1992) 851–858. [40] Z.H. Chohan, H.A. Shad, J. Enz. Inhib. Med. Chem. 23 (2008) 369–379. [41] Z.H. Chohan, H. Pervez, A. Rauf, K.M. Khan, C.T. Supuran, J. Enz. Inhib. Med. Chem. 19 (2004) 417–423. [42] Z.H. Chohan, C.T. Supuran, Appl. Organomet. Chem. 19 (2005) 1207–1214. [43] G.G. Mohamed, Z.H. Abd El-Wahab, Spectrochim. Acta A 61 (2005) 1059–1068.