Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

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Supramolecular spectroscopic and thermal studies of azodye complexes A.Z. El-Sonbati ⇑, M.A. Diab, A.A. El-Bindary, Sh.M. Morgan 1 Chemistry Department, Faculty of Science, Damietta University, Damietta, 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

 The ligands (HLn) and their Cu(II)

complexes were prepared and characterized.  The molecular and electronic structures of the investigated compounds (HLn) have been discussed.  ESR calculations support the characterization of the structures of the complexes geometries.  Thermal data suggested that the ligands are more stable as compared to complexes.  Different thermodynamic parameters were discussed.

a r t i c l e

i n f o

Article history: Received 21 December 2013 Received in revised form 5 February 2014 Accepted 14 February 2014 Available online 25 February 2014 Keywords: Azodye rhodanine derivatives Cu(II) supramolecular structure ESR study TGA and DSC Thermodynamic parameters

R

R N N H2O

C S

Cu

NH O

OAc

HN S C S

O

N

Cu N

C S NH O

OH2

a b s t r a c t A series of heterocyclic ligand of copper(II) complexes have been synthesized by the reaction of copper(II) acetate with 5-(40 -derivatives phenylazo)-2-thioxothiazolidin-4-one (HLn) yields 1:1 and 1:2 (M:L) complexes depending on the reaction conditions. The elemental analysis, spectral (IR and ESR), conductance, magnetic measurements, and thermogravimetric analysis (TGA) are used to characterize the isolated complexes. It is found that the change of substituent affects the thermal properties of azodye rhodanine derivatives and their Cu(II) complexes. The molecular and electronic structures of the investigated compounds (HLn) were also studied using quantum chemical calculations. According to intramolecular hydrogen bond leads to increasing of the complexes stability. The data revealed that the coordination geometry around Cu(II) in all complexes (1–4) exhibit a trans square planar by NO monobasic bidentate and the two monobasic bidentate in octahedral complexes (5–7). Electronic, magnetic data and ESR spectra proposed the square planar structure for all complexes (1–4) under investigation. The value of covalency factor ðb1 Þ2 and orbital reduction factor K accounts for the covalent nature of the complexes. The activation thermodynamic parameters, such as activation energy (Ea), enthalpy (DH), entropy (DS), and Gibbs free energy change of the decomposition (DG) are calculated using Coats–Redfern and Horowitz–Metzger methods. Ó 2014 Elsevier B.V. All rights reserved.

Heterocyclic azodye attracted a considerable interest and play an important role in development of the chemistry of chelates ⇑ Corresponding author. Tel.: +201060081581; fax: +20 572403868. E-mail address: [email protected] (A.Z. El-Sonbati). Abstracted from her Ph.D.

http://dx.doi.org/10.1016/j.saa.2014.02.037 1386-1425/Ó 2014 Elsevier B.V. All rights reserved.

.2H2O

S

N

R

Introduction

1

H 2O N

S

system due to their applicability as potential ligands for a large number of metal ions. Diab et al. and his school survey reveal an excellent work devoted to synthesis and characterization of azodyes as well as their complexes [1–8]. It is well known that azodyes and their metal complexes have been widely used in different fields such as dying of textile fibers [9], biological studies [10], and high technology areas [11]. The variety of coordination modes of rhodanine derivatives has been demonstrated in a number of

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

complexes and their biological applications have been of considerable interest [12]. Although, neither structural chemistry nor coordinating studies have been reported on ligands containing both azo and rhodanine function groups. Data from our laboratory [13–17] have demonstrated that the bidentate azodyes ligands play a key role in making new complexes with transition metal ions. However, little is known concerning the constitution of these complexes, as well as the chemistry involved in their preparation, or the structural and coordination in such complexes. It has been shown from the IR spectral data [13–20] that the hydrogen bonding plays an important role in biological systems. Moreover, Jorgensen and El-Sonbati et al. [21,13–19] found out that the stability of multiple hydrogen bonded depends not only on the number of hydrogen bonds but also on the hydrogen bonding pattern. The importance of clarifying the structure and stability of hydrogen-bonded complexes has opened up an area of surface science that has attracted a considerable attention in the environmental chemistry. The azo group can act as a proton acceptor in hydrogen bonds [14,19,22–24]. The role of hydrogen bonding in azo aggregation has been accepted for sometime. Intarmolecular hydrogen bonds involving OH group with the AN@NA group increased their stabilities through chelate ring structure [14,19,24,25]. The strength of the hydrogen bond of compounds depends on the nature of substituents present in the coupling component from the aryl azo group. Chelating rings formed by NH  N bonds are less stable than corresponding rings formed by OH  N bonds [26,27]. The objectives of the present work are the synthesis of 5-(40 derivatives phenylazo)-2-thioxo-4-thiazolidinone (HLn) (Fig. 1) and their Cu(II) complexes. The thermogravimetric analysis (TGA) and differential scanning calorimetry analysis (DSC) studies for HLn. Study the molecular and electronic structures of the investigated compounds (HLn). The Cu(II) complexes are subjected to elemental, thermogravimetric analysis, spectral studies (IR and ESR), conductance and magnetic measurements for the purpose of structural elucidation. The optimum conditions and the stoichiometry for the reaction of azodye rhodanine derivatives with Cu(II) in solution were considered. In addition to the activation thermodynamic parameters are calculated using Coats–Redfern and Horowitz– Metzger methods.

Experimental All the chemicals used were of British Drug House (BDH) quality.

Synthesis of 5-(40 -derivatives phenylazo)-2-thioxothiazolidin-4-one (HLn) In a typical preparation, 25 ml of distilled water containing 0.01 mol hydrochloric acid were added to aniline (0.01 mol) or pderivatives. To the resulting mixture stirred and cooled to 0 °C, a solution of 0.01 mol sodium nitrite in 20 ml of water was added dropwise. The formed diazonium chloride was consecutively coupled with an alkaline solution of 0.01 mol 2-thioxo-4-thiazolidinone, in 10 ml of pyridine as shown in Scheme 1. The colored precipitate, which formed immediately, was filtered through sintered glass crucible, washed several times with water and ether. The crude products was purified by recrystallization from hot ethanol, yield  65% then dried in vacuum desiccator over P2O5. The ligands were also characterized by elemental analysis (Table 1), 1 H NMR and IR spectroscopy.

311

General method of synthesis of complexes A hot ethanolic solution containing the azodyes (HLn) was mixed with a hot ethanolic solution of Cu(OAc)2H2O (1 mmol). The mixture was then refluxed on a water bath for 10 h and allowed to cool whereby the solid complexes were separated, which filtered off, washed several times with ethanol, dried and kept in a desiccator over dried CaCl2. The analytical data are given in Table 2. Measurements Elemental microanalyses of the separated ligands and solid chelates for C, H, and N were performed in the Microanalytical Center, Cairo University, Egypt. The analyses were repeated twice to check the accuracy of the analyzed data. The metal content in the complexes was estimated by standard methods [28]. X-ray diffraction analyses of the powder HL2 and its complex [Cu(L2)(OAc)(OH2)]2H2O (4) were performed at room temperature by a Philips Xray diffractometer equipped with utilized monochromatic Cu Ka radiation (k = 1.5418 Å). The X-ray tube voltage and current were 40 kV and 30 mA, respectively. The 1H NMR spectrum was obtained with a JEOL FX90 Fourier transform spectrometer with DMSO-d6 as the solvent and TMS as an internal reference. The infrared spectra were recorded as KBr discs using a Perkin–Elmer 1340 spectrophotometer. Ultraviolet–Visible (UV–Vis) spectra of the compounds were recorded in Nuzol solution using a Unicom SP 8800 spectrophotometer. The magnetic moment of the prepared solid complexes was determined at room temperature using the Gouy’s method. Mercury(II) (tetrathiocyanato)cobalt(II), [Hg{Co(SCN)4}], was used for the calibration of the Gouy’s tubes. Diamagnetic corrections were calculated from the values given by Selwood [29] and Pascal’s constants. Magnetic moments were calculated using the equation, leff: ¼ 2:84½T vcoor: 1=2 . ThermograviM metric analysis (TGA) measurements were investigated using Simultaneous Thermal Analyzer (STA) 6000 (Central Laboratory, Tanta University, Egypt) with scan rate 15 °C/min under dynamic nitrogen atmosphere in the temperature range from 50 to 800 °C. ESR measurements of powdered samples were recorded at room temperature (Central Laboratory, Tanta University, Egypt) using an X-band spectrometer utilizing a 100 kHz magnetic field modulation with diphenyl picrylhydrazyle (DPPH) as a reference material. The conductance measurement was achieved using Sargent Welch scientific Co., Skokie, IL, USA. The molecular structures of the investigated compounds were optimized initially with PM3 semiempirical method so as to speed up the calculations. The resulting optimized structures were fully re-optimized using an initio Hartree–Fock (HF) [30] with 6-31G basis set. The molecules were built with the Gauss View 3.09 and optimized using Gaussian 03W program [31]. The corresponding geometries were optimized without any geometry constraints for full geometry optimizations. Frequency calculation was executed successfully, and no imaginary frequency was found, indicating minimal energy structures. Quantum chemical parameters such as the highest occupied molecular orbital energy (EHOMO), the lowest unoccupied molecular orbital energy (ELUMO), energy gap (DE), dipole moment (l) and the total energy for the investigated molecules were calculated. Results and discussion Structure of the ligand All five ligands HL1–HL5, gave satisfactory elemental analysis (Table 1). The molecular structures of these ligands are such that

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O HN S

OH

C

C

HN

CH S

H

O

C

HN

O HN

C

H

C N

N N N

(A)

S

C

C

N N

S

R

S

C

C S

C

C S

R

(C)

R

(B)

S

N

N

(D)

R

R

H O

OH HN S

C

C

C S

or N N

N N

R

C

C S

HN

S

N N

OH

NH

C

C

C

C C

S

C

S

S

C

N N

S

NH

HO

R R

(E)

R

or S

S

N

HN

O HN

(F)

C

C

S

C

C

C S

NH

C

NH

O

N

R

(G) n=1 n=2 n=3 n=4 n=5

R = OCH3 R = CH3 R=H R = Cl R = NO2

Fig. 1. Structure of azodye rhodanine derivatives (HLn).

they can exist in three tautomeric forms as shown in Fig. 1. Detailed solution and solid states studies of these ligands were carried out to establish their geometry. As shown in Table 1, the values of yield% and/or melting point is related to the nature of the p-substituent as they increase according to the following order p-(NO2 > Cl > H > CH3 > OCH3) [16–19]. This can be attributed to the fact that the effective charge experienced by the d-electrons increased due to the electron withdrawing p-substituent (HL4 and HL5) while it decreased by the electrons donating character of (HL1 and HL2). The X-ray diffraction, XRD, patterns of the as-synthesized HL2 and [Cu(L2)(OAc)(OH2)]2H2O (4) powder are shown in Fig. 2. Many peaks are observed which indicate the polycrystalline nature of the as-synthesized HL2 ligand and [Cu(L2)(OAc)(OH2)]2H2O (4). Geometrical structures and electronic properties of the investigated compounds and their protonated forms were calculated by optimizing their bond lengths, bond angles and dihedral angles. The calculated molecular structures with the optimized bond lengths are shown in Fig. 3.

According to the frontier molecular orbital theory, FMO, the chemical reactivity is a function of interaction between HOMO and LUMO levels of the reacting species [32]. The EHOMO often associated with the electron donating ability of the molecule to donate electrons to appropriated acceptor molecules with low-energy, empty molecular orbital. Similarly, ELUMO indicates the ability of the molecule to accept electrons. The lower value of ELUMO indicates the high ability of the molecule is to accept electrons [33,34]. While, the higher is the value of EHOMO of the compound, the easier is its offering electrons. The HOMO and LUMO are shown in Fig. 4. Quantum chemical parameters of organic compounds are obtained from calculations such as energies of the highest occupied molecular orbital, EHOMO, energies of the lowest unoccupied molecular orbital ELUMO, total, binding and electronic energies; heats of formation and dipole moments as presented in Table 3. Additional parameters such as separation energies, DE, absolute Electronegativities, v, chemical potentials, Pi, absolute hardness, g, absolute softness, r, global electrophilicity, x [35], global softness, S, and additional electronic charge, DNmax, have been calculated according to the following equations [36]:

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R

NaNO2 HCl

NH2

N Cl

+

R

0-5 oC

N

Diazonium Salt O

alkaline solution 0-5 oC S R

HN C S

R

R HO

S

HN N

NH

O

S

N N

C S

C S O

NH

Keto form

Hydrazon form

=

S

N N

C S

HO

N

NH

Enol form

n=1, R = OCH3 (HL1); n=2, CH3 (HL2); n=3, H (HL3); n=4, Cl (HL4); and n=5, NO2 (HL5) Scheme 1. The formation mechanism of azodye rhodanine derivatives (HLn). Table 1 Analytical data of azo rhodanine derivatives. a

Compound

Empirical formula

Yield%

M.p.°C

HL1

C10H9N3O2S2 Red C10H9N3OS2 Dark Orange C9H7N3OS2 Pale Yellow C9H6N3OS2Cl Light Orange C9H6N4O3S2 Dark Yellow

37.45

221

47.81

231

42.19

237

51.37

248

66.09

245

HL2 HL3 HL4 HL5

Calc. (Exp.)% C

H

N

44.93 (44.82) 47.79 (47.88) 45.55 (45.68) 39.78 (39.65) 38.29 (38.42)

3.39 (3.25) 3.61 (3.76) 2.97 (2.80) 2.23 (2.35) 2.14 (2.25)

15.72 (15.85) 16.72 (16.61) 17.71 (17.85) 15.46 (15.58) 19.85 (19.98)

a

The analytical data agree satisfactory with the expected formulae represented as given in structures HL1–HL5. Air-stable, colored, insoluble in water, but soluble in hot ethanol, and soluble in coordinating solvent.

DE ¼ ELUMO  EHOMO v ¼ ðEHOMO2þ ELUMO Þ

g ¼ ELUMO 2 EHOMO r ¼ g1 ; Pi ¼ v S ¼ 21g ;

2

x ¼ Pi2g

DNmax ¼  Pig

The concepts of the parameters v and Pi are related to each other. The inverse of the global hardness is designated as the softness r [37]. From the obtained data (Table 3) we can deduced that: Absolute hardness g and softness r are important properties to measure the molecular stability and reactivity. A hard molecule has a large energy gap and a soft molecule has a small energy gap. Soft molecules are more reactive than hard ones because they could easily offer electrons to an acceptor.

1

H NMR spectra

The 1H NMR spectra of azodye rhodanine derivatives are in agreement with the proposed structures. Signal for CH ( 4.42 ppm), favoring formation of an intramolecular hydrogen bond with the AN@NA (azodye) group. Electron-withdrawing substituents reduce the intramolecular hydrogen bond as indicated by the marked shift of the hydroxyl signal to higher field in the p-NO2 and p-Cl compounds. Electron-donating substituents give the opposite effect, arising from the increasing basicity of the azonitrogen. The broad signals assigned to the OH protons at 11.36–11.88 ppm are not affected by dilution. The previous two protons disappear in the presence of D2O. Absence of ACH proton signal of the ligand moiety indicated the existence of the ligand in the azo-enol form. According to El-Sonbati et al. [14–19,25],

Table 2 Analytical and magnetic moments of Cu(II) complexes. Complex

[Cu(L1)(OAc)(OH2)]2H2O (1) [Cu(L3)(OAc)(OH2)]2H2O (2) [Cu(L5)(OAc)(OH2)]2H2O (3) [Cu(L2)(OAc)(OH2)]2H2O (4) [Cu(L1)2(OH2)2] (5) [Cu(L3)2(OH2)2] (6) [Cu(L5)2(OH2)2] (7)

leff. B.M. 1.83 1.85 1.87 1.84 1.92 1.96 2.04

Calc. (Exp.)% C

H

N

S

M

32.54 (32.47) 33.76 (33.60) 32.00 (31.86) 28.51 (28.40) 38.00 (37.84) 37.79 (37.62) 32.65 (32.44)

3.84 (3.72) 3.99 (3.82) 3.64 (3.44) 3.02 (2.89) 3.17 (3.00) 2.80 (2.64) 2.12 (2.01)

9.49 (9.13) 9.85 (9.46) 10.18 (9.87) 9.07 (8.78) 13.30 (13.04) 14.70 (14.49) 16.93 (16.66)

14.46 (14.07) 15.01 (14.86) 15.51 (15.27) 13.82 (13.65) 20.27 (19.74) 22.40 (22.03) 19.35 (19.07)

14.36 (14.06) 14.90 (14.72) 15.40 (15.17) 13.72 (13.47) 10.06 (9.87) 11.12 (10.84) 9.61 (9.48)

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(1) Intramolecular hydrogen bond between the nitrogen atom of the AN@NA system and hydrogen atom of the hydroxy hydrogen atom (Fig. 1C). This is evident by the presence of a broad band centered at 3460 cm1. (2) Hydrogen bonding of the OH  N type between the hydroxy hydrogen atom and the N-ph group (Fig. 1C). (3) Intermolecular hydrogen bonding is possible forming cyclic dimer through NH  O@C (G), OH  N@N (F) or OH  OH (E). (Fig. 1).

3000

Intensity, (A.U)

2400

1800

(b)

1200

600

(a) 0 10

20

30

40

50

60

70

80

2θ (Degree) Fig. 2. X-ray diffraction pattern for (a) HL2 and (b) [Cu(L2)(OAc)(OH2)]2H2O powder form.

The presence of broad band located at 3200 cm1 is strong indication by vNH (Fig. 1D). In general, the low frequency of such region from its normal position is, again due to hydrogen bond property gathered with keto , enol tautomerism. Again the three bands located at 1380, 1340 and 1310 cm1 identified as dOH gathered with the two bands at 1240 cm1 assigned as vCAO are strong indication to keto,enol equilibria. The presence of a medium band at 1605 cm1 assigned to vC@N illustrates the tracing of keto structure (Fig. 1D).

Structure of the metal complexes hydrogen bonding leads to a large deshielding of the protons. The shifts are in the sequence: p-(NO2 > Cl > H > OCH3 > CH3). In the meantime, the 1H NMR of the HL2/HL1 exhibits signals at d(ppm) [3.3(s, 3H, CH3)]/[3.9(s, 3H, OCH3)]. The aromatic protons have resonance at 7.10–7.45 ppm for the ligands. The position of the other proton signals has also been observed in the expected regions and has been shifted only slightly due to the coordination of the ligand to metal ion. Absence of CH proton signal of the rhodanine azo moiety indicated the existence of the ligand in the azo-enol form.

Diazonium coupling of the rhodanine with p-derivatives aromatic amine readily gives rise to azodye ligands. Their reactions with copper(II) ion afford mono/binuclear metal complexes in which the microanalytical data as well as metal and acetate estimations are in good agreement with proposed stoichiometries, Table 2. These complexes are found to be insoluble in common organic solvents but soluble in coordinated solvents. The elemental analysis confirms that the complexes are mono/binuclear and may be formulated as: [Cu(Ln)(OAc)(OH2)]2H2O/[Cu(Ln)2(OH2)2]. The reactions follow:

Infrared spectra of ligands (HLn)

CuðOAcÞ2  H2 O þ HL15 ! ½CuðL15 ÞðOAcÞðOH2 Þ  2H2 O

The infrared spectra of ligands (HLn) give two bands at 3200– 3040 cm1 due to asymmetric and symmetric stretching vibrations of NAH group and intramolecular hydrogen bonding NH  O systems (Fig. 1D), respectively. When the OH group (Fig. 1C) is involved in intramolecular hydrogen bond, the O  N and N  O bond distances are the same. But, if such mechanism is happened in case of intermolecular hydrogen bond, the O  O and O  N bond distances differ. The broad absorption band located at 3400 cm1 is assigned to vOH. The low frequency bands indicate that the hydroxy hydrogen atom is involved in keto , enol (A , B) tautomerism through hydrogen bonding (Fig. 1C). Bellamy [38] made detailed studies on some carbonyl compounds containing ANHA group. The DvNH values were used to study the phenomena of association. On the other hand, the OH group (Fig. 1-B) exhibits more than one absorption band. The two bands located at 1330 and 1370 cm1 are assigned to in-plane deformation and that at 1130 cm1 is due vCAOH. However, the 860 cm1 band is probably due to the out-ofplane deformation of the AOH group. On the other hand, the two bands located at 650 and 670 cm1 are identified as vC@O and NH. Similar to the other investigated compounds, the different modes of vibrations of CAH and CAC band are identified by the presence of characteristic bands in the low frequency side of the spectrum in 600–200 cm1. The infrared spectra of ligands shows medium broad band located at 3460 cm1 due the stretching vibration of some sort of hydrogen of hydrogen bonding. El-Sonbati et al. [18,24] made detailed studies for the different types of hydrogen bonding which are favorable to exist in the molecule under investigation:

þ CH3 COOH CuðOAcÞ2  H2 O þ 2HL15 ! ½CuðL15 Þ2 ðOH2 Þ2  þ 2CH3 COOH where L1–5 = deprotonated HL1, HL2, HL3, HL4, HL5. The composition coordination mode and geometry of the complexes were established on the basis of spectral analyses, conductivity measurements and magnetic properties and are discussed in detail in the following sections.

Conductivity measurements Conductivity measurements have been used in structural elucidation of metal complexes (mode of coordination) within the limits of their solubility. The molar conductance of 103 M of solutions of the complexes in DMSO is measured at room temperature. It is concluded from the results that Cu(II) chelates with HLn ligand under investigation were found to have molar conductance values in the range from 1.85 to 0.55 O1 mol1 cm2 for 1:1 and 1:2 (M:L) complexes, indicating non-electrolytic nature of these compounds and there is no counter ion present outside the coordination sphere of copper complexes [24,39]. This is in accordance with the fact that conductivity values for a non-electrolyte are below 50 O1 mol1 cm2 in DMSO solution [39]. Such a non-zero molar conductance value for each of the complex in the present study is most probably due to the strong donor capacity of DMSO, which may lead to the displacement of anionic ligand and change of electrolyte type [24,39].

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

HL1

HL2

HL3

HL4

HL5 Fig. 3. The calculated molecular structures of the investigated compounds (HLn).

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HOMO

LUMO

HL1

HL2

HL3

HL4

HL5

Fig. 4. The Highest Occupied Molecular Orbital (HOMO) and the Lowest Unoccupied Molecular Orbital (LUMO) of the investigated compounds (HLn).

Table 3 The calculated quantum chemical parameters of the investigated compounds (HLn). Compound

HOMO (a.u)

LUMO (a.u)

DE (a.u)

l (D)

T.E (a.u)

v

g

r

Pi

S

x

DNmax

HL1 HL2 HL3 HL4 HL5

0.327 0.337 0.348 0.355 0.367

0.048 0.047 0.044 0.034 0.004

0.375 0.384 0.432 0.389 0.363

5.358 4.554 3.950 2.001 3.712

1491.580 1416.775 1377.752 1836.627 1581.099

0.139 0.145 0.152 0.161 0.186

0.188 0.192 0.196 0.195 0.182

5.333 5.208 5.102 5.141 5.509

0.139 0.145 0.152 0.161 0.186

2.667 2.604 2.551 2.571 2.755

0.052 0.055 0.059 0.066 0.095

0.744 0.755 0.776 0.825 1.022

Infrared spectra of complexes and nature of coordination The bonding of the metal ion to the ligand can be clarified by comparing the IR-spectra of the complexes with those of the ligands.

The broad/strong absorption bands located at 3450–3434 and at 1710–1695 cm1 region assigned to the t(NH) stretching vibrations mode and carbonyl stretching vibrations. The three bands in the 1600–1500 cm1 region are characteristic for most six-membered aromatic ring system. The frequencies for the

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N@N stretching lie in the region 1440–1435 cm1. The region between 1500–1900 cm1 is due CAN stretching, NAH in plane or out of plane bending and out-of-plane CAH bending vibrations [3–6]. The symmetric and antisymmetric (C@C) stretching vibration modes are expected to exist in this region. Thus, this ligand (HLn) contains four potential donor sites (Fig. 1): (i) the ring nitrogen NH, (ii) the ring CS, (iii) the carbonyl group, and (iv) the nitrogen of azo (N@N) group. However, considering the planarity of the ligand, it is unlikely that this ligand could be tetradentate on a single metal center. Hence, this ligand is expected to be bidentate and the three favorable possibilities of donor sites are: (i) the carbonyl oxygen and (ii) the nitrogen atom of azo (N@N) group or (i) the carbonyl oxygen and (ii) the CS. The coordination of the ring nitrogen (NH) is unlikely due to the Zwitterions [4] formation, thereby lowering the electron density on N. The ring thion (CS) in this ligand is found to be inert towards coordination to copper as revealed by the appearance of the t(C@S)(ring) mode at 820 cm1 of the uncoordinated ligand after complexation. The absence of CH2 and NH2 stretching in the spectra of the ligands (HLn), related to active methane and amino groups respectively, indicating occurrence of azodye diazonium and formation of the proposed skeleton of ligands. In the IR spectra of all metal complexes (1–7) a number of changes are observed:

(4)

(5)

(6)

(7) (1) The appearance of a new bands around 3380 cm1 and two sharp bands at 715 and 420 cm1, the latter two can be assigned to the wagging and rocking modes of vibration of the water molecule, respectively, [39] in the prepared complexes (1–7) may be taken as a strong evidence for the presence of coordinated water. Such region, however, is not initially present in the free ligands. This is confirmed by the elemental analysis of these complexes (Table 2). (2) The N@N stretching frequency of the azo groups is shifted to lower frequency by 15–25 cm1 due to the involvement of one of the azo nitrogen atoms in coordination with metal ion [4,24,40]. This lowering of frequency can be explained by the transfer of electrons from nitrogen atom to the Cu(II) ion due to coordination. (3) In solution and in the presence of Cu(II) ions these compounds exist in a tautomerism equilibrium [24,25] A , B , C (Fig. 1). The tautomeric form (C) react with metal ions by loss of phenolic proton as mononegative chelating agents producing of CO/OH mode of the free ligands. The internally hydrogen enolic band disappeared in the spectra the metal complexes, indicating the deprotonation and formation of metal–oxygen bond. This is further supported by the shifting of t(CAO) towards higher frequency as

(8)

compared to the free ligand due to the conversion of hydrogen bonded structure into a covalent metal bonded structure [41]. Furthermore, the bands in the regions 545–560 and 420– 428 cm1 can be assigned to the stretching modes of the metal to ligand bonds, t(CuAO) and t(CuAN) for 1:1 and 1:2 (M:L) complexes, respectively, [4,17,41]. Absence of t(MAS) band in the far IR spectra gives added evidence for the non-participation of ring sulfur atom in bond formation. The variation in the spectral bands of the CAC, C@C, CAN and CAH different modes of vibrations of the complexes lead to that, probably, the aromaticity of the complex is differ from one to other [14,22]. The band due to vOH is masked by the intense absorption due to water molecule, and hence no definite conclusion can be drawn for this region. However, dOH at 1299 cm1, vCAOH at 1207 and cOH at 822 cm1 display a sharp decrease in their intensities to such an extent those they nearby vanish. This can be taken as an indication for the complete removal of OH-group by the Cu(II) ion reacting with the ligands. Since the ligand reacts with Cu(II) ion, as gathered from the results of elemental analysis, then each Cu(II) ion would displace only one proton from the OH group contained in each ligands molecule. To confirm the presence of coordinating H2O molecules in the complex we carried out thermogravimetric analysis of all the complexes. This study shows loss of weight corresponding to one water molecule in the temperature up to 190 °C, indicating that the water molecule in these complexes is coordinated to the metal ion. The participation of the OH group in Cu(II) complexes is confirmed by the appearance of new bands in Cu(II) complexes is confirmed by the appearance of new bands at 555– 525 cm1 for 1:1 and 1:2 (M:L) complexes respectively, related to the CuAO vibration [24]. The acetate complexes show two new bands at 1600 and 1390 cm1 attributed to tas and ts of the acetate group. The difference between the two bands indicates the monodentate nature of the acetate group [5,8].

On the basis of all these data, the molecular structure of the Cu(II) complexes could be suggested based on: (i) the presence of anion, (ii) the disappearance of C@O, (iii) the coordination of azo-group and (iv) the presence of water. According to the structure shown in Fig. 5 the ligand (HLn) takes its usual anionic form (Ln) to chelate Cu(II) through NA of azo group with enol group (Fig. 1C) as the potential binding sites.

R

R N N H2O

H 2O N

S C S

Cu

NH O

OAc

.2H2O

HN S C S

O

N

S

N

NH

Cu N

C S O

OH2

R

(A)

317

(B)

Fig. 5. Structures of (A) [Cu(Ln)(OAc)(OH2)]2H2O and (B) [Cu(Ln)2(OH2)2].

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14070–14290 cm1 assigned to 2B1g ? 2A1g transition as reported for square planar Cu(II) complexes [44].

Table 4 ESR spectral parameters for the Cu(II) complexes (1–3).

a b

Complexa

gll

g?

gav.

Allb

G

a2

b2

1 2 3

2.210 2.230 2.250

2.040 2.050 2.060

2.097 2.110 2.123

56.57 65.83 76.17

5.614 4.804 4.293

0.421 0.471 0.524

0.80 0.89 0.95

Numbers as given in Table 2. A values in 104 (cm1).

Magnetic and spectral studies Copper(II) complexes (1–4) Magnetic measurements and electronic spectra were conducted in order to obtain information about the geometry of the complexes. The magnetic susceptibility values (1.83–1.87 B.M.) which was consistent with presence of a single unpaired electron [42]. This behavior suggest square planar geometry for the copper(II) complexes [43]. The electronic spectra of [CuLn(OAc)(OH2)]2H2O complexes exhibit a broad band with a maximum at

Copper(II) complexes (5–7) Where spin–spin coupling between unpaired electrons belonging to different copper ions is absent, leff varies between 1.92 and 2.04 B.M., depending on the geometries of the complexes due to difference in orbital contribution [45]. The observed leff values for the Cu(II) complexes in this study are in the 1.92–2.04 B.M. range, corresponding to one unpaired electron. Electronic absorption spectra arise from the electronic transitions within a molecule or ion from a lower to a higher electronic energy level. The transition metal ions generally show a number of d–d transition bands depending on their electronic configuration from d1 to d9 in UV–Vis regions. The copper(II) complexes generally show the d–d transition bands in the range 13100–15500 (dz2 ? dx2y2) (t1), 18640–19100 (dxy ? dx2y2)(t2) and 24410–27330 cm1 (dxz,dyz ? dx2y2)(t3) transitions. The spectra are typical of Cu(II) complexes with an elongated tetragonal. The spectra of all the complexes have been assigned to D4h symmetry with a dx2y2

Table 5 ESR spectral and calculated bonding parameters for the Cu(II) complexes (5–7). gll

g?

gav.

Allb

G

a2

b2

b21

K 2ll

K 2?

K2

d2

5 6 7

2.1839 2.2402 2.2421

2.0216 2.0311 2.0319

2.0757 2.1013 2.1020

39.65 46.72 35.14

6.306 8.238 8.101

0.34 0.42 0.39

0.53 0.58 0.60

1.5 1.6 1.7

0.51 0.67 0.69

0.18 0.24 0.23

0.29 0.38 0.38

0.64 0.77 0.87

Numbers as given in Table 2. A values in 104 (cm1).

80

(a)

2.25

(b)

3

3

75

All x 10-4 (cm -1)

ll

2.24 2

2.23

g 2.22

70 2

65

60

1

2.21

1

-0.4

-0.2

0.0

0.2

0.4

0.6

55 -0.4

0.8

0.60

-0.2

0.0

0.2

0.4

0.6

(c) 7

0.58 6

0.56

0.54 5

0.52 -0.4

-0.2

0.8

Hammett's substituent coefficients (σ R)

Hammett's substituent coefficients (σ R)

β2

a b

Complexa

0.0

0.2

0.4

0.6

0.8

Hammett's substituent coefficients (σ R) Fig. 6. The relation between Hammett’s substituent coefficients (rR) vs. (a) gll, (b) All  104 (cm1) and (c) b2.

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A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

(a)

ESR spectra of copper(II) complexes

(b) g

3

ll

2.25

β

2

0.96

3 0.93

2.24

0.87

ll

2

0.84

2.22

0.81

1

2.21

1 0.78 55

60

65

70

75

80

All x 10 -4 (cm-1 ) Fig. 7. The relation between All  104 (cm1) vs. (a) gll and (b) b2.

ground state. The most active vibration in this point group appears to be b1u symmetry and its efficiency may arise from its being the only out-of-the-xy-plane vibration. The complexes are with one electron sequence i.e. dx2y2 > dz2 > dxy > dxz,dyz [45].

(c)

(a)

(b)

0.96

80

3

2.125

2.115

2

β 0.87

-1

0.90

-4

2.120

g av.

0.93

All x 10 (cm )

3

3

75

2.110

2 70

2 65

0.84

All x 10

2

2.105

-4

-1

(cm )

g av.

2.100

60

1

0.81

β

1 2.095

0.78

2

1

55

0.42

0.44

0.46

0.48

α

0.50

0.52

0.54

2

Fig. 8. The relation between a2 vs. (a) All  104 (cm1), (b) gav. and (c) b2.

2.250

(a)

6

2.032

(b)

7

7

2.235

6

2.030 2.028

2.220

g

ll

g

T 2.026

2.205 2.024 2.190

2.022 5

5

2.020

2.175 0.52

0.56

0.60

0.64

0.68

0.18

k2 ll

Fig. 9. The relation between (a) K 2ll vs. gll and (b) K 2? vs. g ? .

0.20

0.22

k2 T

g

2.23

β2

0.90

2

ESR spectra of copper(II) complexes (1–3) To obtain further information about the stereochemistry and the site of the metal ligand bonding and to determine the magnetic interaction in the metal complexes. ESR spectra of the complexes were recorded in the solid state. The spin Hamiltonian parameters of the complexes were calculated and summarized in Table 4. The analysis of spectra of all copper(II) complexes agree well with the values reported for distorted square planar geometry around Cu(II) ion [46,47]. The spectra shows the relation g ll > g? > ge which is typical of axially symmetric d9 Cu(II) having one unpaired electron in dx2y2 orbital [48]. The absence of signal corresponding to (Ms = ±2) in the half field indicates the absence of any CuACu interaction thus ruling out possibility of dimeric structure. In the axial spectra, the g-values are related with exchange interaction coupling constant (G). The geometric parameter G i.e. the measurement of exchange between the copper centers in the polycrystalline compounds, is calculated by using the expression: G = (gll2.0023)/(g\2.0023). In the present work, G value comes out to be 4.0 which again indicate absence of CuACu interaction, thus supporting proposed monomeric structure. Kivelson and Neiman have reported that gll values less than 2.3 indicates considerable

0.24

TGA

0

Weight loss (%)

DSC 80

100

60

200

40

300

20

400

0

500

-20

600 0

Weight loss (%)

100

(b)

TGA DSC

80

100 200

60

300

40

400

20 0

500

-20

600

(c)

100

TGA

-120

DSC

Weight loss (%)

Heat Flow Endo Down (mW)

(a)

100

Heat Flow Endo Down (mW)

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

80

0

60

120

40

240

20

360

0

480

Heat Flow Endo Down (mW)

320

600

-20 0

100

200

300

400

500

600

700

800

Temperature (°C) Fig. 10. TGA and DSC thermo graphs for (a) HL1, (b) HL3 and (c) HL5.

covalent character of MAL bond and greater than 2.3 indicates ionic character. The present value comes out to less than 2.3 which indicate considerable covalent character of CuAL bond [49]. As a measure of the covalence of the in-plane r-bonding, a2 indicates complete covalent character, whereas a2 = 0.5 denotes 100% covalent bonding, with the assumption of negligible small values of the overlap integral [49]. The b2, the larger the covalence of the bonding. From this analysis, the in-plane bonds of the complexes correspondent to signals 1, 2 and 3 are largely covalent, with the covalent character increasing from 1 to 3. The values of a2 indicate that approximately 90% of the spin population is in

the copper dx2y2 orbital of most of the Cu(II) species concerned. The g values obtained from ESR spectra indicate, as has been pointed out by Kivelson and Neiman [49], that covalent bonding reduces the magnitude of the g factor. ESR spectra of copper(II) complexes (5–7) The ESR technique is very sensitive tool for obtaining the formation on the chemical environments and the coordination states of Cu(II) in the materials. The intensity of the ESR signals from a paramagnetic Cu(II) ion dependents on both the condition spheres of this ion and the degree of the dispersion [50]. The spin Hamiltonian, orbital reduction and bonding parameters of these complexes are given in Table 5. The observed gll values for all complexes are less than 2.3 commensurate a significant covalent character of the metal–ligand bond in agreement with the observation of Kivelson and Neiman [49]. The trend gll > g\ > ge > 2.0023 for these complexes suggests that the unpaired electron is localized in the dx2y2 orbital [51] of the Cu(II) ion. The g values reflect that the Cu(II) center has a tetragonal distorted octahedral geometry with dx2y2 orbital as a ground state [24,52]. The gll values are an important function for indicating covalent character of MAL; for ionic character, gll > 2.3 and for covalent character gll < 2.3. In the present complex, the gll is less than 2.3 indicating appreciable covalent character for the CuAL bond. The unpaired electron in this 3d9 case assigned to 3dx2y2 orbital, and the overlapping of this antibonding orbital with the ligand 2s and 2p r orbital is often determined by the use of the following equation:

a2 ¼ All =Po þ ðg ll  2:0023Þ þ K o ðg ?  2:0023Þ þ 0:04 where a2 gives an approximate indication of the strength of the interaction between the metal and the ligands. Po is the dipolar contribution to the hyperfine splitting value A, which is negative quantity, and usually assigned the free ion value of 0.036 cm1. The constant term of Ko = 3/7 is equated to the Fermi hyperfine contact term of the free ion, which corrects for the Fermi contact contributions from excited state configurations of Cu(II), notably the 3s1 3d10 and 3s2 3d8 4s1 configurations [53]. This is usually considered to be a constant term for Cu(II) complexes since there is little ligand orbital density at the copper nucleus and the ratio of copper s to d character is assumed to be unchanged in the presence of ligands [54]. The addition of 0.04 to the sum is an approximate correction due to the molecular orbital coefficients of the complex that arise from in plan p bonding (dxy ; b21 ) and out of plan p bonding (dxz, dyz, b2). Thus a2 represents the extent to which the unpaired electron resides on the central metal ion in the dx2y2 orbital (B1g ground state) and reflects the extent of r-bonding to the ligand. Consequently, a decline in the value of a2 indicates an increase in the covalencey of the bond. The in-plane r-covalence parameter a2 values account for the fraction of the unpaired electron density to be populated around the Cu(II) ion. In addition, the g values are related to the G-factor by the expression, G = (gll2)/ (g?  2) = 4, which measures the exchange interaction between copper centers in the solid state. According to Hathaway [55], if the value of G is greater than 4, the exchange interaction between Cu(II) centers in the solid state is negligible, whereas when it is less

Table 6 Weight losses percentage of HL1, HL3 and HL5. Compounds

HL1 HL3 HL5

First stage

Second stage

Third stage

Remaining weight after 650 °C

Temperature (°C)

Weight loss (%)

Temperature (°C)

Weight loss (%)

Temperature (°C)

Weight loss (%)

200 145 112

61.9 59.8 5.4

320 400 225

38.1 39.9 33.6

– – 284

– – 58.7

Ash Ash Ash

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A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328 Table 7 Thermal analyses data of the Cu(II) complexes (1–4).

a

Complexa

Temp. range (°C)

Found mass loss (calc.)%

Assignment

(1)

45–189 189–331 331–506 506–709

7.24 16.67 24.43 14.38

(8.13) (17.39) (24.17) (14.46)

Loss of two water molecule in outside of the coordination sphere Loss of one coordinated water and one coordinated acetate group Further decomposition of a part of the ligand (C7H7O) Evolution of SO2 gas leaving CuO residue with contaminated carbon atoms

(2)

45–178 178–273 273–514 514–560 560–640

7.91 18.61 18.06 6.64 15.43

(8.73) (18.66) (18.42) (6.79) (15.51)

Loss of two water molecule in outside of the coordination sphere Loss of one coordinated water and one coordinated acetate group Further decomposition of a part of the ligand (C6H4) Evolution of N2 gas Evolution of SO2 gas leaving CuO residue with contaminated carbon atoms

(3)

45–189 189–288 288–400 400–592 592–670

4.15 16.72 10.46 13.72 22.44

(7.87) (16.83) (10.05) (13.98) (22.73)

Loss of two water molecule in outside of the coordination sphere Loss of one coordinated water and one coordinated acetate group Evolution of NO2 gas Evolution of SO2 gas Decomposition of a part of the ligand (C6H4N2) leaving CuO residue with contaminated carbon atoms

(4)

40–189 189–316 316–533 533–670

5.22 18.00 20.47 15.45

(8.44) (18.05) (21.33) (15.00)

Loss of two water molecule in outside of the coordination sphere Loss of one coordinated water and one coordinated acetate group Further decomposition of a part of the ligand (C7H7) Evolution of SO2 gas and CuO residue with contaminated carbon atoms

Numbers as given in Table 2.

(a)

80

60

40

20 0

100

200

300

400

500

600

700

(b)

100

Weight loss (%)

Weight loss (%)

100

80

60

40

20

800

0

100

200

Temperature (°C)

400

500

600

700

800

Temperature (°C)

(c)

100

80

60

40

(d)

100

Weight loss (%)

Weight loss (%)

300

80

60

40 20 0

100

200

300

400

500

600

700

800

0

100

200

300

400

500

600

700

800

Temperature (°C)

Temperature (°C)

Fig. 11. TGA curves for complexes (a) [Cu(L1)(OAc)(OH2)]2H2O, (b) [Cu(L3)(OAc)(OH2)]2H2O, (c) [Cu(L5)(OAc)(OH2)]2H2O and (d) [Cu(L2)(OAc)(OH2)]2H2O.

than 4,a considerable exchange interaction exists in the solid complex. The calculated G value for these complexes indicates that there are no interactions between the copper centers (see Table 5). The a2 value of 0.5 indicates complete covalent bonding, while that of 1.0 suggests complete ionic bonding. The out-of-plane pbonding (d2) and in-plane p-bonding (b) parameters are calculated using the following equations:

b2 ¼ ðg ll  2:0023ÞE=  8ka2 2

2

d ¼ ðg ?  2:0023ÞE=  2ka

Here k = 828 cm1 for free Cu(II) ion and E is the electronic transition energy. This is also confirmed by orbital reduction factors K, which can be estimated using the following equations:

K ll ¼ a2 b2 K ? ¼ a2 d2 Significant information about the nature of bonding in the Cu(II) complex can be derived from the relative magnitudes of Kll and K ? . In the case of pure r-bonding, Kll K ? = 0.77, whereas Kll < K ?

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HL1

-12.5

HL1

-11.5

ln [-ln (1- α) / T2 ]

ln [-ln (1- α) / T2 ]

-13.0

-12.0

-12.5

-13.0

-13.5 -14.0 -14.5 -15.0 -15.5

-13.5

o

211-346 C

498-624 oC

-16.0

0.00161 0.00168 0.00175 0.00182

0.00111 0.00114 0.00117 0.00120 0.00123 0.00126 0.00129

0.00189 0.00196

(1/T(K))

(1/T(K))

-11.5 -12.0

HL3

-12.0

HL3

-12.5

ln [-ln (1- α) / T2 ]

ln [-ln (1- α) / T2]

-12.5 -13.0 -13.5 -14.0 -14.5

-13.0 -13.5 -14.0 -14.5 -15.0

-15.0 -15.5

502-643 oC

-15.5

158-401 oC

-16.0 0.0016 0.0017 0.0018 0.0019 0.0020 0.0021 0.0022 0.002

0.00108 0.00112 0.00116 0.00120 0.00124 0.00128

(1/T(K))

(1/T(K))

HL 5

-12.0

-13.0

ln [-ln (1- α ) / T 2 ]

-12.8

ln [-ln (1- α) / T2 ]

HL5

-12.5

-13.6 -14.4 -15.2 -16.0

-13.5 -14.0 -14.5 -15.0

-16.8 -17.6

-15.5

225-271 oC 0.00186

0.00189

0.00192

0.00195

0.00198

-16.0 0.00110

443-630oC 0.00115

(1/T(K))

0.00120

0.00125

0.00130

0.00135

(1/T(K)) Fig. 12. Coats–Redfern (CR) of the ligands (HLn).

implies considerable in-plane p-bonding, while for out-of-plane p-bonding Kll > K ? . Molecular orbital coefficients a2 (in-plane r-bonding), b2 (in-plane p-bonding) and d2 (out-plane p-bonding) were calculated using the last equations (see Table 5). The lower value of b2 indicates that the in-plane p-bonding is more covalent than the in-plane r-bonding, the data in a good agreement with the data reported earlier [24]. The a2 values for copper(II) complexes indicates a considerable covalencey in the bonding between the Cu(II) ion and the metal. In this study, a2 is less than b2 indicating that in-plane r-bonding is more covalent than in-plane p-bonding. These ESR data showed that:  gll, All  104 cm1 and b2 values are dependent of the substituents effect of the p-position of the ligand and can be order as: p(OCH3 < H < NO2) (Fig. 6). The reverse order is observed for the electronic spectral data.

 The value of All  104 cm1 increases with increasing gll and b2 (Fig. 7).  The value of a2 increases with increasing All  104 (cm1), gav and b2 (Fig. 8).  K 2ll and K 2? increase with increasing gll and g ? (Fig. 9). Thermal analyses Thermogravimetric analysis of ligands (HLn) The thermal properties of ligands were characterized on the basis of thermogravimetric analysis (TGA) and differential scanning calorimetry (DSC) methods in the temperature range 50–800 °C. The TGA and DSC curves for HLn (where n = 1, 3 and 5) is shown in Fig. 10. The temperature intervals and the percentage of loss of masses are listed in Table 6. For HL1 and HL3 there are two steps of the loss of masses, while for HL5 there are three steps. Fig. 10(a)

323

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328 -11.5

[Cu(L 1 )(OAc)(OH 2 )].2H 2 O

-12.0

-12.0

(1)

[Cu(L1 )(OAc)(OH2 )].2H2 O (1)

-12.5

ln [-ln (1- α) / T 2 ]

ln [-ln (1- α) / T2 ]

-12.5 -13.0 -13.5 -14.0

-13.0 -13.5 -14.0 -14.5

-14.5 -15.0

142-240 oC

-15.0 0.00196

0.00203

0.00210

0.00217

0.00224

0.00231

301-418 oC

-15.5 0.00145

0.00238

0.00150

0.00155

0.00165

0.00170

0.00175

-11.5

-11.5

[Cu(L 3)(OAc)(OH 2)].2H 2O

[Cu(L 3 )(OAc)(OH 2 )].2H 2 O

-12.0

(2)

(2)

-12.0

-12.5

ln [-ln (1- α) / T2 ]

ln [-ln (1- α) / T2 ]

0.00160

(1/T(K))

(1/T(K))

-12.5

-13.0

-13.5

-13.0 -13.5 -14.0 -14.5

o

127-372 C

496-715 oC

-15.0

-14.0 0.0016

0.0017

0.0018

0.0019

0.0020

0.0021

0.0022

0.00100

0.00105

0.00110

(1/T(K))

0.00115

0.00120

-12.0

-11.5

[Cu(L 5)(OAc)(OH 2)].2H 2O

-12.0

[Cu(L 5 )(OAc)(OH 2 )].2H 2 O

-12.5

(3)

(3)

ln [-ln (1- α) / T 2 ]

-12.5

ln [-ln (1- α) / T 2 ]

0.00125

(1/T(K))

-13.0 -13.5 -14.0 -14.5

-13.0 -13.5 -14.0 -14.5

-15.0 -15.5

-15.0

o

o

189-292 C

349-403 C -15.5

-16.0 0.00180 0.00185 0.00190 0.00195 0.00200 0.00205 0.00210 0.00215

0.00150

0.00152

0.00154

0.00156

0.00158

0.00160

(1/T(K))

(1/T(K)) -12.0

[Cu(L 2)(OAc)(OH 2)].2H 2O

-12.0

[Cu(L 2)(OAc)(OH 2)].2H 2O

(4)

ln [-ln (1- α) / T 2 ]

ln [-ln (1- α) / T 2 ]

(4)

-12.5

-12.5

-13.0

-13.5

-14.0

-13.0

-13.5

-14.0

-14.5

o

o

599-715 C

142-399 C -14.5

-15.0 0.00150

0.00165

0.00180

0.00195

0.00210

0.00225

0.001008

0.001036

0.001064

0.001092

0.001120

0.001148

(1/T(K))

(1/T(K)) Fig. 13. Coats–Redfern (CR) of the Cu(II) complexes (1–4).

reveals that the HL1 decomposes in two steps, the first stage, weight loss starts at 200 °C with a weight loss percentage of

61.9%. The second stage, weight loss starts at 320 °C with a weight loss percentage of 38.1%. The HL3 decomposes in two steps as

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A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328 0.2

0.4

HL 1

-0.2 0.0

log [log (Wα/W γ)]

log [log (Wα /W γ )]

HL 1

0.0

0.2

-0.2 -0.4 -0.6

-0.4 -0.6 -0.8 -1.0 -1.2

-0.8

211-346 oC

498-624 oC

-1.4

-1.0 -60 -50 -40 -30 -20 -10

0

10

20

30

40

50

60

-1.6

70

-60 -50 -40 -30 -20 -10

θ (K)

10

20

30

40

50

60

70

θ (K) 0.3

HL 3

0.0

0

HL 3

0.0

log [log (Wα /W γ)]

log [log (Wα /W γ)]

-0.3 -0.3

-0.6

-0.6

-0.9

-0.9

-1.2

-1.2

-1.5 o

158-401 C

-1.8 -120 -100

-80

-60

-40

-20

0

20

40

502-643 oC

-1.5 -60

60

-40

-20

0

20

40

60

80

θ (K)

θ (K) 0.2

HL 5

0.0

0.0

HL 5

-0.3 -0.2

log [log (Wα /W γ)]

log [log (Wα /W γ)]

-0.6 -0.9 -1.2 -1.5 -1.8 -2.1

-0.4 -0.6 -0.8 -1.0 -1.2

-2.4

-1.4

o

225-271 C

-2.7

443-630 oC

-1.6

-25

-20

-15

-10

-5

0

5

10

15

20

-80

-60

-40

-20

0

20

40

60

80

θ (K)

θ (K) Fig. 14. Horowitz–Metzger (HM) of the ligands (HLn).

shown in Fig. 10(b), the first stage, weight loss starts at 145 °C with a weight loss percentage of 59.8%. The second stage, weight loss starts at 400 °C with a weight loss percentage of 39.9%. Fig. 10(c) reveals that the HL5 decomposes in three steps, the first stage, weight loss starts at 112 °C with a weight loss percentage of 5.4%. The second stage, weight loss starts at 225 °C with a weight loss percentage of 33.6%. The third stage, weight loss starts at 284 °C with a weight loss percentage of 58.7%. The DSC data for HL3 and HL5 ligands show two exothermic and endothermic peaks except HL1 ligand show one exothermic peak is shown in Fig. 10(a–c). For the HL1 ligand the recorded DSC curve show only one exothermic peak at 220 °C. Simultaneously, the recorded DSC curve of the HL3 ligand show two exothermic and endothermic peaks at 190 °C and 180 °C, respectively, while the recorded DSC curve of the HL5 ligand show two exothermic and endothermic peaks at 245 °C and 120 °C, respectively. These

endothermic peaks can be attributed to the loss of residual solvent trapped in the ligands matrices. Thermogravimetric analysis of complexes (1–4) The TGA of the isolated complexes was taken as a proof for the existing of water molecules as well as the anions in the coordination sphere. The thermal analyses data for the [CuLn(OAc)(OH2)]2H2O complexes (n = 1, 2, 3, and 5) are summarized in Table 7. It can be seen that the TGA curves of the complexes [CuLn(OAc)(OH2)]2H2O complexes show loss of masses down to 100 °C, indicating that the presence of two water molecules in outside of the coordination sphere, while up to 189 °C above which the dehydration begins are shown in Fig. 11. In the temperature range 189–331 °C, the lost of the one coordinated water molecule as well as the lost of acetate group. The weight loss stages (270–500 °C) are due to the decomposition of a part of the ligand.

325

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328 0.2

0.0

[Cu(L 1)(OAc)(OH 2)].2H 2O

-0.2

0.0

[Cu(L 1)(OAc)(OH 2)].2H 2O (1)

(1) log [log (Wα /W γ)]

log [log (Wα /W γ)]

-0.2

-0.4 -0.6 -0.8 -1.0 -1.2

-0.4 -0.6 -0.8 -1.0 -1.2

-1.4

142-240 oC

-1.6 -50

-40

-30

-20

-10

0

10

20

30

301-418 oC

-1.4 -60

40

-50

-40

-30

-20

θ (K)

-10

0

θ (K)

10

20

30

40

50

0.4 0.2

[Cu(L 3)(OAc)(OH 2)].2H 2O

0.6

[Cu(L 3)(OAc)(OH 2)].2H 2O

(2)

0.4

(2)

0.2

-0.2

log [log (Wα /Wγ)]

log [log (Wα /W γ)]

0.0

-0.4 -0.6 -0.8 -1.0

0.0 -0.2 -0.4 -0.6 -0.8

-1.2

-1.0

o

-1.4

127-372 C

-1.6 -120 -100 -80 -60 -40 -20

0

20

40

60

80 100 120

496-715 oC

-1.2 -100 -80

-60

-40

-20

0.0

0.0

[Cu(L 5)(OAc)(OH 2)].2H 2O

-0.2

0

20

40

60

80

100

θ (K)

θ (K)

(3)

[Cu(L 5)(OAc)(OH 2)].2H 2O (3)

-0.2

-0.4

log [log (Wα /W γ)]

log [log (Wα /Wγ)]

-0.4

-0.6 -0.8

-0.6

-1.0

-0.8

-1.2

-1.0

-1.4 -1.6

-1.2

-1.8

189-292 oC

349-403 oC

-1.4

-2.0 -50

-40

-30

-20

-10

0

10

20

30

40

-25

-20

-15

-10

0.2 0.0

0

5

10

15

20

0.4

[Cu(L 2)(OAc)(OH 2)].2H 2O

[Cu(L 2)(OAc)(OH 2)].2H 2O (4)

(4)

0.2

log [log (Wα /Wγ)]

-0.2

log [log (Wα /W γ)]

-5

θ (K)

θ (K)

-0.4 -0.6

0.0

-0.2

-0.8 -1.0

-0.4

-1.2 -1.4

142-399 oC

599-715 oC

-0.6

-1.6 -100 -80 -60 -40 -20

0

20

40

60

80 100 120

-60

-50

-40

-30

-20

θ (K) Fig. 15. Horowitz–Metzger (HM) of the Cu(II) complexes (1–4).

-10

0

θ (K)

10

20

30

40

50

326

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

Table 8 Thermodynamic data of the thermal decomposition of HLn and their Cu(II) complexes (1–4). Compounda

Decomposition temperature (°C)

Method

Parameter Ea (kJ mol1)

HL1

211–346 498–624

HL3

158–401 502–643

HL5

225–271 443–630

1

142–240 301–418

2

127–372 496–715

3

189–292 349–403

4

142–399 599–715

a

CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM CR HM

45.8 53.6 180 193 46.9 48.4 162 173 329 338 108 120 70.9 80.8 102 114 31.2 39.2 108 119 106 115 235 250 31.8 41.0 171 187

r A (s1) 1

7.58  10 4.18  102 4.43  108 6.40  109 7.11  101 1.17  102 2.34  107 2.50  108 2.54  1030 1.70  1032 1.51  104 2.05  105 5.53  105 1.40  107 1.31  106 2.10  107 3.31  100 3.59  101 1.59  104 5.78  104 3.74  108 7.73  109 9.36  1016 2.09  1018 2.10  100 3.60  101 4.27  107 2.01  108

DG (kJ mol1)

214 200 87.9 65.8 215 210 113 92.2 333 367 173 152 139 112 134 111 240 220 173 163 85.3 60.0 73.5 99.3 244 220 108 95.4

41.2 49 174 186 42.3 43.8 155 166 325 333 101 113 67.1 76.9 96.5 108 26.9 34.9 101 112 102 111 230 244 27.3 36.5 163 179

159 159 247 241 161 160 250 245 151 142 242 236 131 129 181 179 152 150 254 255 146 142 182 180 160 156 264 268

0.99221 0.99398 0.99355 0.99169 0.99322 0.99788 0.98442 0.99156 0.99453 0.99452 0.99432 0.9939 0.99122 0.98595 0.99104 0.99064 0.99523 0.98744 0.98712 0.98669 0.99596 0.99460 0.99590 0.99435 0.99161 0.98130 0.98947 0.99112

The thermodynamic activation parameters of decomposition processes of complexes namely activation energy (Ea), enthalpy (DH), entropy (DS), and Gibbs free energy change of the decomposition (DG) are evaluated graphically by employing the Coast– Redfern [58] and Horowitz–Metzger [59] methods. Coast–Redfern equation The Coast–Redfern equation, which is a typical integral method, can represent as: a

dx A ¼ ð1  aÞn u

Z

T2

T1

  Ea dt exp  RT

ð1Þ

For convenience of integration, the lower limit T1 usually taken as zero. This equation on integration gives:

    lnð1  aÞ Ea AR ¼  ln  þ ln RT uEa T2

ð2Þ

A plot of left-hand side (LHS) against 1/T was drawn (Figs. 12 and 13). Ea is the energy of activation in J mol1 and calculated from the slop and A in (s1) from the intercept value. The entropy of activation DS in (J K1 mol1) calculated by using the equation:

   Ah R DS ¼ 2:303 log kB T s

where kB is the Boltzmann constant, h is the Plank’s constant and Ts is the TG peak temperature. Horowitz–Metzger equation The Horowitz–Metzger equation is an illustrative of the approximation methods. These authors derived the relation:

" # 1  ð1  aÞ1n Eh ¼ log ; 1n 2:303RT 2s

Calculation of activation thermodynamic parameters

0

DH (kJ mol1)

Numbers as given in Table 2.

According to literature the azo bonds in the azo metal complexes breakdown when the temperature is higher than 260 °C [25,56,57]. The final weight losses (500–670 °C) are largely attributed to complete decomposition of other part of the ligand molecule with evolution of SO2 gas. The remaining final product is CuO.

Z

DS (J mol1 K1)

ð3Þ

for n – 1

ð4Þ

when n = 1, the LHS of Eq. (4) would be log[log(1a)] (Figs. 14 and 15). For a first order kinetic process, the Horowitz–Metzger equation may write in the form:

   Wa Eh ¼ log log  log 2:303 Wc 2:303RT 2s

ð5Þ

where h = TTs, wc = waw, wa = mass loss at the completion reaction; w = mass loss up to time t. The plot of log [log (wa/wc)] vs. h was drawn and found to be linear from the slope of which Ea was calculated. The pre-exponential factor, A, calculated from equation:

Ea RT 2s

¼h

A  i u exp  RTEas

ð6Þ

The entropy of activation, DS, is calculated from Eq. (3). The enthalpy activation, DH, and Gibbs free energy, DG, calculated from:

DH ¼ Ea  RT

ð7Þ

DG ¼ DH  T DS

ð8Þ

The calculated values of Ea, A, DS, DH and DG for the decomposition steps for ligands (HLn) and their complexes (1–4) are summarized in Table 8. The complex (3) is the highest value of Ea. This can be attributed to the fact that the effective charge experienced by the d-electrons

A.Z. El-Sonbati et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 127 (2014) 310–328

increases due to the electron withdrawing p-substituent NO2 while it decreases by the electron donating character of OCH3 and CH3. This indicates that, the complex (3) is more thermally stable than the other complexes. From the values of the energy of activation (Ea) of the ligands (HLn) and their complexes (1–4), it is observed that the complexes (1–4) are less stable than the ligand (HLn) and may be attributed to the loss of water molecules from complexes. These results are inconsistent with our previous results [60]. From these results we can say that, the high values of activation energies reflect the thermal stability of the compounds. The entropy of activation energies reflects the thermal stability of the compounds. The negative values of activation entropies (DS) indicate a more ordered activated compounds than the reactants and/or the reactions are slow [61]. The values of DG is positive considered as favorable or spontaneous reaction. Conclusions

[9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21]

In this work, synthesis and characterization of a series of 5-(40 derivatives phenylazo)-2-thioxothiazolidin-4-one (HLn) and their copper(II) complexes (1–7) are reported. The analytical and physicochemical analyses confirmed the composition and the structure of the newly obtained compounds. The results obtained can be summarized as follows: 1. Elemental analysis, IR and molar conductivity data are used to proof the stoichiometry and formulation of the complexes. A square planar and octahedral geometries are assumed for all complexes, based on the magnetic data, spectral (ESR and visible) and thermal studies. 2. The new compound behaves as monobasic bidentate ligand when react with Cu(II) salt and undergo coordination through azodye nitrogen and enolic oxygen atom. 3. ESR calculations support the characterization of the structures of the complexes geometries. 4. The data revealed that the coordination geometry around Cu(II) in all complexes (1–4) exhibit a trans square planar by NO monobasic bidentate and the two monobasic bidentate in octahedral complexes (5–7). 5. The molecular and electronic structures of the investigated compounds (HLn) have been discussed. 6. The thermal properties of the ligands (HLn) and their Cu(II) complexes were investigated by thermogravimetry analysis (TGA) and different thermodynamic parameters are calculated using Coats–Redfern and Horowitz–Metzger methods. 7. Thermal data suggested that the ligands are more stable as compared to complexes.

[22]

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