Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 125 (2014) 61–66

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Spectrophotometric investigation of interaction between iodine and pentadentate Schiff base ligands Z. Khouba a,b, T. Benabdallah a, U. Maschke b,⇑ a Laboratoire de Chimie et d’Electrochimie des Complexes Métalliques (LCECM), Université des Sciences et de la Technologie d’Oran-Mohamed Boudiaf, BP 1505, El-Mnaouer 31000, Algérie b Unité Matériaux et Transformations (UMET), UMR CNRS N° 8207, Université Lille1 – Sciences et Technologies, 59655 Villeneuve d’Ascq Cedex, France

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

 Structural effect of Schiff bases on the

stability constant of iodine complexes.  Spectrophotometric evaluation of ionization potentials of selected ligands.  Thermodynamic parameters well described by the modified Benesi– Hildebrand equation.

a r t i c l e

i n f o

Article history: Received 7 June 2013 Received in revised form 20 December 2013 Accepted 8 January 2014 Available online 21 January 2014 This publication is dedicated to the memory of Professor Ali Hassoune Al-Taiar from USTO, who initiated the work presented here. He passed away suddenly on February 27, 2013. We will always keep him in thankful and honourable memory. Keywords: Pentadentate Schiff base Charge transfer complex Benesi–Hildebrand Interaction Ionization potential Spectrophotometry

a b s t r a c t The interaction between iodine as an electron acceptor (A), and three pentadentate Schiff bases, 1,3bis(salicylideneamino)-2-propanol (SB1), 1,3-bis(2-hydroxy-1-naphthylideneamino)-2-propanol (SB2), and 1,3-bis[1-(pyridine-2-yl)methylideneamino]-2-propanol (SB3), as electron donor systems (D), was studied spectrophotometrically in methanol at 28 °C. Equilibrium constants KAD and molar extinction coefficients eAD of the donor–acceptor complexes (AD) were determined using the modified Benesi–Hildebrand equation in conjunction with the non linear fit analysis. The method shows the formation of 1:1 type complexes as major species in solution. The free energy changes DG° and the energy of the charge transfer band ECT were also calculated for all complexes. The iodine complex derived from SB2 seems to be more stable than those derived from SB3 and SB1. On the other hand, the ionization potential ID of each Schiff base was estimated from the corresponding complex band energy, using an empirical equation. An inverse relationship between ID and KAD values was found. Blue and red shift observed for the 445 nm band of iodine were also discussed on the basis of theoretical considerations. Ó 2014 Elsevier B.V. All rights reserved.

Introduction

⇑ Corresponding author. Tel.: +33 3 20 33 63 81. E-mail address: [email protected] (U. Maschke). http://dx.doi.org/10.1016/j.saa.2014.01.044 1386-1425/Ó 2014 Elsevier B.V. All rights reserved.

The interactions between electron-donors and electron-acceptors occur in different states of matter. These interactions can be strong, medium or weak, depending on the nature of the donor

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and the acceptor systems. Thus, stable inorganic complexes usually result from the interactions between strong base ligands and high acid metal ions. Interactions might also exist between organic entities, some playing the role of donors and the others the role of acceptors, forming organic–organic complexes. The charge transfer from electron-donors (D) to electron-acceptors (A) represents an interaction leading to the formation of charge transfer complex (CTC), known for their low reaction enthalpies (in the order of a few kJ), and high rates of formation or decomposition, so that the reaction appears to be spontaneous [1]. Several authors have studied CTC, especially those derived from p-donors, such as polycyclic aromatic compounds, or n-donors such as aliphatic amines, amine oxides, ethers, picolines, pyrazolones as well as virus DNA nucleotide, with r-acceptors such as iodine and bromine, or p-acceptors such as trinitrobenzene, benzoquinones, etc. . .[2–8]. It is known that the spectrophotometric study in solution, of interactions between iodine and various donor systems has attracted much attention of researchers for over half a century [9–13]. This type of interaction was revealed for the first time by Benesi and Hildebrand, when they observed a new absorption band of iodine-benzene and iodine-mesitylene mixtures in neutral solvents, like carbon tetrachloride and n-heptane, due to neither electron donor nor acceptor. Such phenomenon was considered as CTC formation [14]. Although extensive studies have been carried out on CTC resulting from iodine and different donor molecules [15–21], there are still some organic systems whose interactions with iodine have not been intensively studied. Among them, Schiff bases seem to be of particular interest, because of their basicity and stability, in particular those derived from salicylaldehyde and derivatives [22,23]. Indeed, interactions of Schiff bases towards various metal ions have been extensively investigated both in solution and solid state. Those derived from salicylaldehyde or derivatives are the most exploited in terms of complexation, and showed different behavior depending on the environment and the nature of the complexing agent. Such systems have been evaluated for their biological activity, their catalytic effect as well as for their liquid–liquid extraction properties [24–27]. The present investigation was undertaken to describe the spectrophotometric behaviour of some selected pentadentate Schiff bases as n- or p-donors, towards molecular iodine as r-acceptor. These Schiff bases, presented in Fig. 1, have the same alkyl chain and differ from their aromatic nucleus, influencing thus their electronic effect. The presence of different coordination sites of these systems, such as imine and alkoxo groups, phenolic and naphtylic hydroxides, leads to a large variety of interactions towards acceptor systems. To the best of our knowledge, there are only a few reports of this type of Schiff bases in the literature [28–31]. Experimental part

H

C

N

N

C

N

C

N

C

H

OH OH HO

SB1

H

C

N

H

OH OH HO

SB2

H

C

N N

OH

H

N

SB3 Fig. 1. Chemical structures of the pentadentate Schiff bases.

Nuclear magnetic resonance The 1H NMR and 13C NMR spectra were taken on an FT-NMR (300 MHz) Bruker instrument. Deuterated chloroform (with 0.03% tetramethylsilane) and deuterated dimethylsulfoxide (with 0.06% tetramethylsilane) were used as NMR solvents. They were obtained in high quality (>99%), from Eurisotop (Gif sur Yvette, France), and used as received.

Elemental analysis The elemental analysis was realized by the central service of analysis of the CNRS. The following elements were analysed: C, H, N, O.

Ultraviolet/visible spectroscopy The UV–visible measurements were taken on a double beam Perkin Elmer spectrophotometer Lambda 20, using Hellma quartz cells of 1.0 cm path length.

Materials and techniques Sample preparation The bi-sublimated iodine (>99%) was purchased from Merck, stored in dessiccator and used without further treatment. Methanol was from Merck and was used as received. The pentadentate Schiff bases SB1, SB2 and SB3 (Fig. 1), were prepared and purified according to the general procedure described in the literature [32–34]. The purity of the ligands was checked by using Fourier transform infrared spectroscopy, nuclear magnetic resonance, elemental analysis and UV–visible spectroscopy. Fourier transform infrared spectroscopy Fourier transform infrared FTIR were recorded in the transmission mode using a Perkin Elmer 2000 model. The number of accumulated scans was 16 with a spectral resolution of 4 cm1.

The initial methanolic solutions of iodine and Schiff bases were freshly prepared, in 50 ml volumetric flasks, using the following concentrations: [I2]0 = 103 M, for [SB1-I2] and [SB3-I2] complexes, [I2]0 = 5  4 10 M, for the [SB2-I2] complex. [SB1]0 = [SB3]0 = 2  103 M, and [SB2]0 = 5  104 M. Then, a series of nine samples was prepared for each complex, in 10 ml volumetric flasks, maintaining a fixed volume of iodine (1 ml), and varying quantity of Schiff bases (from 1 to 9 ml). The volume of all solutions was adjusted to 10 ml by methanol. Then, the spectra of all complexes were recorded at room temperature (28 °C) between 200 and 600 nm.

Z. Khouba et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 125 (2014) 61–66

Results and discussion Characterization of the Schiff bases Infrared spectra of the Schiff bases show principally strong sharp bands situated in the range from 1630 to 1647 cm1, characteristic for C@N stretching vibrations. The broad absorption bands observed in the region 3170–3510 cm1 can be assigned to the stretching vibrations of aliphatic and aromatic hydroxyl groups. The 1H NMR data are represented by the multiplet signals at 6.7–8 ppm, corresponding to the aromatic protons. Absorption signals due to methylene protons (ACH2A) occur in the range between 3.4 and 4.2 ppm. The aliphatic (HOACHA) proton is characterized by a multiplet in the range from 4.3 to 5.6 ppm. The azomethine protons (ACH@NA) exhibit a chemical shift in the range from 8.3 to 9 ppm. The 13C NMR spectra show the signals in the range from 106 to 161 ppm, corresponding to the aromatic carbons: six signals for SB1, ten for SB2 and five for SB3. The carbon of the azomethine group exhibits a chemical shift at 167, 177 and 163 ppm for SB1, SB2, and SB3, respectively. The methylene carbons display signals at 63.2, 54.7 and 65 ppm, and the CHOH carbons are represented by the signals situated at 70.4, 68.8 and 69.6 ppm, respectively. The C, H, N contents of the Schiff bases are in accordance with their chemical formulae: SB1 (C17H18N2O3); C, 68.45% (68.57%); H, 6.04% (6.0%), N, 9.40% (9.46%). SB2 (C25H22N2O3); C, 75.37% (75.28%); H, 5.52% (5.49%); N, 7.03% (7.12%). SB3 (C15H16N4O); C, 67.16% (67.13%); H, 5.97% (6.02%); N, 20.89% (20.96%) (the calculated values are shown in brackets).

AD = eD[D]0l: initial absorbance of the donor in the absence of the acceptor. eA and eD represent the molar extinction coefficients of the acceptor and the donor respectively, at kCT, and l is the optical path length. [A]0 and [D]0 are the initial concentrations of A and D in the mixture, respectively. In general, for complexes of the type 1:1 (one donor to one acceptor), the variation of [A]0/Aa against 1/[D]0 leads to a linear relationship, with a slope defined by 1/(KADea) and the intercept 1/ea. A rough estimation of eAD and KAD of the complexes can be deduced, and used for the nonlinear fit analysis [43]. As mentioned previously, both acceptor and donor absorb at a maximum of the wavelength kCT of the complex AD. The absorbance of the mixture, A, can be written as follows:

A ¼ eA ½Al þ eD ½Dl þ eAD ½ADl

½AD ¼ K AD ½A0 ½D0 =ð1 þ K AD ½D0 Þ ðA  AD Þ=AA ¼ ð1=K þ c½D0 Þ=ð1=K þ ½D0 Þ where c = (e

AD

D

–e )/e . Thus, Eq. (7) becomes the following form:



y ¼ ða þ b xÞ=ða þ xÞ

ð2Þ

where Aa (Eq. (3)) and ea (Eq. (4)) represent the apparent absorbance and the apparent extinction coefficient of the resulting complex AD, respectively:

Aa ¼ A  AA  AD

ð3Þ

ea ¼ eAD  eA  eD

ð4Þ

A, represents the absorbance of the mixture at kCT. AA = eA[A]0l: initial absorbance of the acceptor in the absence of the donor.

ð8Þ

a and b are the fitting parameters and can be expressed as follows:

b ¼ c ¼ ðeAD  eD Þ=eA

½A0 =Aa ¼ 1=ðK AD ea ½D0 Þ þ 1=ea

ð7Þ

A

Previous studies have been reported in the literature using the Benesi–Hildebrand Eq. (1), or one of its derivatives, to investigate interactions between some donors and acceptors in different solvents [35–38]. Some thermodynamic parameters can be established for 1:1 type complexes derived from a donor D and an acceptor A, if the initial concentration of the donor in the mixture, [D]0, is greater than that of the acceptor, [A]0, and neither D nor A absorb at a maximum of the absorbing band kCT of the resulting charge transfer complex AD:

where Abs, is the absorbance of the complex at a maximum of its charge transfer band, kCT. KAD and eAD are respectively the association constant, and the molar extinction coefficient of the complex AD. Several authors have modified this equation if the complex is not the only absorbing species in the mixture as it is the case in the present study [39–42]. In this paper, we are interested in calculating some thermodynamic parameters of three iodine-Schiff base complexes, [SB1-I2], [SB2-I2] and [SB3-I2], using the modified Benesi–Hildebrand Eq. (2), known as Ketelaar equation, in order to take into account a significant absorbance of each component at kCT of the resulting complexes:

ð6Þ

Eq. (7) below is obtained in conjunction with Eqs. (5) and (6):

a ¼ 1=K AD

ð1Þ

ð5Þ

where [AD] represents the concentration of the complex in the mixture and is expressed by the following formula:

Theory and formulae

½A0 =Abs ¼ 1=ðK AD eAD ½D0 Þ þ 1=eAD

63

ð9Þ ð10Þ

where y = (A–AD)/AA (the reduced absorbance at kCT) and x = [D]0 (the initial concentration of the donor). The stability constant KAD and the molar extinction coefficient eAD can be calculated from the parameters a and b obtained by the nonlinear curve fit, by using Eqs. (9) and (10). The Gibbs energy changes of the complex formation, DG°, have been deduced from Eq. (11):

DG ¼ RTLnK AD

ð11Þ

R is the ideal gas constant, and T is the absolute temperature (301 K). The charge transfer energy, ECT, for each complex, was determined by using Eq. (12):

ECT ¼ hmCT ¼ hc=kCT

ð12Þ

where h is the Planck constant, c is the speed of light, and kCT is the wavelength of the charge transfer band. The ionization potentials ID of the three Schiff bases were also established, on the basis of the optical data of their iodine complexes, by using an empirical Equation (13) [44,45].

hmCT ¼ ID  C 1 þ ½ðC 2 =ID  C 1 Þ

ð13Þ

The constants C1 and C2, related to iodine, have the following values:C1 = 5, 45 ± 0, 26 eV and C2 = 1, 63 ± 0, 43 eV2 [46]. Analysis of the obtained UV–visible data The spectrophotometric data of methanolic solutions of the starting reagents I2, SB1, SB2, SB3, and their resulting complexes [SB1-I2], [SB2-I2] and [SB3-I2] are summarized in Table 1. As it can be seen from the data, all complexes absorb between 400 and 460 nm. This region of the spectra is characterized by an interference phenomenon, so each component absorbs significantly at the wavelength, kCT, of the complex. For this reason, the Ketelaar equation (2) was employed to estimate the values of the formation

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Table 1 Wavelengths kmax of SB1, SB2, SB3, I2, and their resulting complexes, [SB1-I2], [SB2-I2] and [SB3-I2], in methanol at 301 K.

a b

Systems

Selected wavelengths kmax (nm)

SB1 SB2 SB3 I2 [SB1-I2] [SB2-I2] [SB3-I2]

267, 230, 245, 225, 246, 250, 236,

277, 268, 253, 288, 253, 310, 251,

315, 308, 278 360, 289, 330, 355,

400 328, 400, 420 445a 360, 418b 400, 420, 456b 420b

The visible band of iodine in methanol. Charge transfer bands of the iodine-Schiff base complexes.

constants KAD and the molar extinction coefficients eAD of all complexes, involving the corrected absorbance of each sample, according to Eq. (3). The absorbance was measured directly at a maximum of the new absorbing band, kCT, of each complex (418, 456 and 420 nm for [SB1-I2], [SB2-I2], and [SB3-I2] respectively). In general, the values of kmax of each of the starting reagents were found to be different from those of their corresponding complexes. In fact, the new absorption bands for [SB1-I2], [SB2-I2] and [SB3-I2] complexes occur at 418, 456, and 420 nm respectively (Table 1). As originally described by Mulliken, these absorptions arise from excitation from the ground state of the complex WN to the excited state WE. For this reason, such systems are generally called CTC [47]. It was observed that all complexes absorb at the lower wavelength region of the spectra (higher energies). This may be a result of the partial transfer of the charge from the donor (Schiff base) to the acceptor (iodine). However, it was reported that the metal complexes derived from these Schiff bases or analogues, ab-

sorb in general at longer wavelengths (lower energies). This is originally due to the complete transfer of the electron from the ligands to the metal ions (LM), forming the stable complexes with defined structures [48–50]. As it can be seen from the data displayed in Table 1, the 445 nm visible band of iodine undergoes a blue shift in the case of [SB1-I2] and [SB3-I2] complexes (by 27 and 25 nm respectively), and a red shift in the case of [SB2-I2] complex (by 11 nm). The mechanism of the ‘‘blue shift’’ of the visible band of iodine has already been discussed by Mulliken and Person [47]. They considered that the excited orbital for the visible band of iodine, ru MO, produces a strong non-bonded repulsion toward donor partner in complex formation, because of its big size. Nagakura has given a similar explanation based on LCMO approximation [51]. According to this method, the charge transfer attraction originates from the interaction between the above mentioned ru MO and the non-bonding orbital on the nitrogen atom of the donor system (triethylamine). As the result of this interaction, the former orbital becomes higher in energy (and the latter lower). Thus the visible band of iodine is shifted towards shorter wavelengths. Probably, the two Schiff bases under consideration SB1 and SB3, interact with iodine molecule in such a manner. The heteroatom can be respectively an oxygen or nitrogen. The red shift observed in the case of [SB2-I2] complex, can be probably explained by the high conjugation of the naphtylic groups of the ligand, causing the bathochromic effect of the absorbing band of the complex. Calculation A plot of [A]0/Aa against 1/[D]0 (Eq. (2)), for each complex, leads to the straight lines displayed in Fig. 2a–c, with good correlation factors (R2 > 0.98). The linear correlation obtained for each system

Fig. 2. Plots of [I2]0/Aa against 1/[SB]0 according to the Ketelaar equation (2), for (a) [SB1-I2], (b) [SB2-I2] and (c) [SB3-I2] complexes. The experimental data are shown by symbols whereas continuous lines represent the theoretical linear fits.

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Table 2 Values of wavelengths kCT, energies of absorption bands ECT, estimated molar extinction coefficients eAD and formation constants KAD, standard free energy changes DG° of iodine-SB complexes, and ionization potentials ID of the three Schiff bases in methanol, at 301 K. Complexes

kCT (nm)

ECT (eV)

eAD (L cm1 mol1)104

KAD (L mol1)103

DG° (kJ mol1)

ID (eV)

[SB1-I2] [SB2-I2] [SB3-I2]

418 456 420

2.967 2.725 2.952

3.078 2.590 0.208

0.237 4.020 0.993

13.68 20.76 17.27

7.69 (SB1) 7.29 (SB2) 7.67 (SB3)

Fig. 3. Reduced absorbance (A–AD)/AA against [SB]0 in methanol, at 301 K. The experimental results were analyzed by a fitting procedure using Eq. (8) at 418 nm for [SB1-I2] (a), at 456 nm for [SB2-I2] (b) and at 420 nm for [SB3-I2] (c). The experimental data are shown by symbols whereas continuous lines represent the theoretical non-linear fits.

Table 3 Values of the fitting parameters a and b, calculated formation constants KAD and molar extinction coefficients eAD of iodine-SB complexes in methanol, at 301 K. Complexes

[SB1-I2] [SB2-I2] [SB3-I2]

Fitting parameters

Calculated values

a (mol L1)10+4

b

KAD (L mol1)103

eAD (L cm1 mol1)104

42 2.5 10

29.4 9.1 2.1

0.263 4.348 1.136

2.8577 2.8086 0.2013

predicts the formation of 1:1 complexes as predominated species in solution, as it was reported by several authors [35–42]. If other species, with different stœchiometries, are present in solution, they could be detected in the plots by the changes of the slope beyond a certain concentration of the donor. The estimated values of KAD and eAD were obtained from the intercept of the line with the ordinate, and from the gradient. The results are summarized in Table 2, and then applied for the nonlinear fit analysis according to Eq. (8). The corresponding plots are displayed in Fig. 3a–c. For all complexes, the values of b parameter are greater than 1.1, and then Person’s range of concentrations remains valid in this study (Table 3) [43]. The calculated values of KAD and eAD deduced

from the nonlinear plots (Table 3) are relatively close to the estimated ones. Such results lead to the conclusion that the assumption made for single 1:1 complexes should be valid. On the other hand, the convergence of all results confirms that the experimental measurements were realized with a minimum of errors and a high reproducibility in each case. The values of KAD constants related to the iodine complexes derived from SB1, SB2, and SB3 are respectively 263, 4348 and 1136 dm3 mol1. From these results, the order of stability of the complexes can be established as follows: [SB2-I2] > [SB3-I2] > [SB1-I2]. SB2 seems therefore to lead to the more stable complex, which can be explained by the high electronic density of the naphtyl clusters, increasing the nucleophily of the ligand, con-

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Z. Khouba et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 125 (2014) 61–66

solidating the ligand-iodine links and stabilizing the complex. The [SB3-I2] complex seems also to be more stable than its analogue [SB1-I2]. This may be explained by the high affinity of the pyridinic nitrogen towards iodine as revealed by Tsubomura [11]. The values of the molar extinction coefficients reflecting the intensities of the charge transfer bands of [SB1-I2], [SB2-I2], and [SB3I2] complexes, are 28577, 28086 and 2013 dm3 cm1 mol1, respectively. The values of the standard Gibbs energy changes DG°, reflecting the spontaneity of the complex formation, are evaluated with the help of Eq. (5), and the results are displayed in Table 2. The values are of the order of a few kJ/mol, indicating a charge transfer character of the complexes [1]. The ionization potentials of the Schiff bases were calculated from Eq. (7), using the values of kCT of each corresponding complex, and the results are summarized in Table 2. An inverse relationship could be established between the ionization potential ID of each Schiff base and the formation constant KAD of its corresponding iodine complex since the stability of the complex is related to the ability of the charge transfer from the donor to the acceptor. As it can be seen from the results, SB2 appears to have the lowest ionization potential and the highest stability constant, so it can be concluded that the charge transfer in this system is easy compared to the other donors (SB3 and SB1). Conclusions The spectrophotometric determination of the initial equilibrium constants KAD and molar extinction coefficients eAD, for three iodine-Schiff base complexes, was applied using the Ketelaar equation. This choice was related to the appreciable absorption of the starting reagents at the maximum of the absorbing charge transfer bands kCT of the complexes. The results revealed the formation of 1:1 type complexes in all cases. The non-linear fit analysis allowed to obtain KAD and eAD values for the same complexes. As a result, a good agreement was found between experimental data and theoretical predictions. This can be explained by the low standard deviation observed in Fig. 3a–c between theoretical and experimental curves. The [SB2-I2] complex absorbs at longer wavelength (456 nm) and seems to be more stable than the other ones. The high electronic density of the naphtyl clusters of SB2 may be the reason of its high stability. The ionization potentials ID of the Schiff bases SB1, SB2 and SB3 were estimated on the basis of their spectrophotometric data, using an empirical equation. This rapid method seems to be interesting since the use of the standards techniques such as the photo-ionization is rather difficult [1,52– 54]. The values of ID are in the range between 7.29 and 7.69 eV. The inverse relationship found between KAD and ID values shows the dependence of the stability of the complexes on the ionicity of their corresponding ligands. Acknowledgements The authors gratefully thank the government of Algeria, the University Lille 1, as well as the CNRS for their financial support. References [1] R. Foster, Organic Charge Transfer Complexes, first ed., Academic Press, London, New York, 1969. pp. 42-44 and 125-139. [2] T. Karuna, K. Neelima, G. Venkateshwarlu, P.Y. Swamy, J. Sci. Ind. Res. 65 (2006) 808–811.

[3] H. Razaq, R. Qureshi, N.K. Janjua, F. Saira, S. Saleemi, Spectrochim. Acta A 70 (2008) 1034–1040. [4] L. Shahada, A. Mostafa, E.-M. Nour, H.S. Bazzi, J. Mol. Struct. 933 (2009) 1–7. [5] J.S. Askar Ali, P.A. Abdullah Mahaboob, M. Nizam Mohideen, J. Chem. Pharm. Res. 4 (2012) 901–912. [6] J. Peters, W.B. Person, J. Am. Chem. Soc. 86 (1964) 10–14. [7] A.A. Ibrahim, Microchim. J. 58 (1998) 175–181; A.A. Ibrahim, Microchim, A. J. P. A. C 5 (2011) 507–514. [8] G.A.-R. Yuldasheva, G.M. Zhidomirov, A.I. Ilin, Nat. Sci. 3 (2011) 573–579. [9] C. Reid, R.S. Mulliken, J. Am. Chem. Soc. 76 (1954) 3869–3874. [10] L.E. Orgel, R.S. Mulliken, J. Am. Chem. Soc. 79 (1957) 4839–4846. [11] H. Tsubomura, J. Am. Chem. Soc. 82 (1959) 40–45. [12] M.J.S. Dewar, A.R. Lepley, J. Am. Chem. Soc. 83 (1961) 4560–4563. [13] W.B. Person, J. Am. Chem. Soc. 87 (1965) 167–170. [14] H.A. Benesi, J.H. Hildebrand, J. Am. Chem. Soc. 71 (1949) 2703–2707. [15] P.A. Clark, T.J. Lerner, S. Hayes, S.G. Fisher, J. Phys. Chem. 80 (1976) 1809–1811. [16] A.H. Al-Taiar, M.A. Al-Ekabe, F.A. Al-Mashadane, M. Hadjel, J. Coord. Chem. 55 (10) (2002) 1143–1146. [17] M.S. Refat, S.M. Taleb, I. Grabchev, Spectrochim. Acta A 61 (2005) 205–211; M.S. Refat, M.M. Ibrahim, A.A.M. Moussa, J. Appl. Spectrosc. 78 (6) (2012) 772– 781. [18] S.A. Mizyed, M. Ashram, R. Saymeh, D. Marji, Z. Naturforsch. 60b (2005) 1133– 1137. [19] T.M. El-Gogary, M.A. Diab, S.F. El-Tantawy, Spectrochim. Acta A 66 (2007) 94– 101. [20] H.S. Bazzi, S.Y. ElQaradawi, A. Mostafa, E. Nour, J. Mol. Struct. 879 (2008) 60– 71. [21] S.I.M. Zayed, Cent. Eur. J. Chem. 7 (2009) 870–875. [22] M. El-Aasser, F. Abdel-Halim, M. Ashraf El-Bayoumi, J. Am. Chem. Soc. 93 (1971) 590–592. [23] A.A.H. Saeed, A.R.J. Al-Azzawy, Spectrosc. Lett. 25 (1992) 777–788. [24] Y.-Y. Ye, H.-D. Xian, J.-F. Lue, G.-L. Zahao, Molecules 14 (2009) 1747–1754. [25] S. Deshpande, D. Srinivas, P. Ratnasamy, J. Catal. 188 (1999) 261–269. [26] M. Hadj Youcef, T. Benabdallah, H. Ilikti, H. Reffas, Solvent Extr. Ion Exch. 26 (2008) 534–555. [27] H. Reffas, T. Benabdallah, M. Hadj Youcef, H. Ilikti, J. Chem. Eng. Data 55 (2010) 912–918. [28] M. Mikuriya, T. Sasaki, A. Anjiki, S. Ikenoue, T. Tokii, Bull. Chem. Soc. Jpn. 65 (1992) 334–339; M. Mikuriya, Y. Hatano, E. Asato, ibid. 70 (1997) 2495–2507. [29] R.J. Butcher, G. Diven, G. Erickson, J. Jasinski, G.M. Mockler, R.Y. Pozdniakov, Ekk Sinn, Inorg. Chim. Acta 239 (1995) 107–116. [30] A. Mukherjee, M. Nethaji, A.R. Chakravarty, Polyhedron 23 (2004) 3081–3085. [31] Z. Khouba, T. Benabdallah, U. Maschke, Mol. Cryst. Liq. Cryst. 502 (2009) 121– 129; Z. Khouba, T. Benabdallah, U. Maschke, Phys. Procedia 2 (2009) 1305–1311. [32] T.G. Campbell, F.L. Urbach, Inorg. Chem. 12 (1973) 1836–1840. [33] U. Casellato, P.A. Vigato, M. Vidali, Coord. Chem. Rev. 23 (1977) 31–117. [34] W. Mazurek, K.J. Berry, K.S. Murray, M.J. O’Connor, M.R. Snow, A.G. Wedd, Inorg. Chem. 21 (1982) 3071–3080. [35] R. Foster, D.L. Hammick, A.A. Wardley, J. Chem. Soc. (1953) 3817–3820. [36] R.P. Lang, J. Am. Chem. Soc. 84 (1962) 1185–1192. [37] K.R. Gorda, D.G. Peiffer, Polymer 32 (1991) 2060–2063. [38] J.K. Kochi, R. Rathore, P. Le Magueres, J. Org. Chem. 65 (2000) 6826–6836. [39] N.J. Rose, R.S. Drago, J. Am. Chem. Soc. 81 (1959) 6138–6141. [40] K.F. Wong, S. Ng, Spectrochim. Acta A 32 (1976) 455–456. [41] B.K. Seal, H. Sil, D.C. Mukherjee, Spectrochim. Acta A 38 (1982) 289–292. [42] T.R. Bera, D. Sen, R. Ghosh, Spectrochim. Acta A 45 (1989) 985–991. [43] G. Sarova, M.-N. Berberan-Santos, J. Phys. Chem. B 108 (2004) 17261–17268. [44] S.H. Hastings, J.L. Franklin, J.C. Schiller, F.A. Matsen, J. Am. Chem. Soc. 75 (1953) 2900–2905. [45] G. Briegleb, J. Czekalla, Z. Elektrochem. 63 (1959) 6–12. [46] R.S. Becker, E. Chen, J. Chem. Phys. 45 (1966) 2403–2410. [47] R.S. Mulliken, W.B. Person, Molecular Complexes, first ed., Wiley-Interscience, New York, London, 1969. [48] Qi-L. Zhang, Bi-X. Zhu, L.F. Lindoy, G. Wie, Inorg. Chem. Commun. 11 (2008) 678–680. [49] A.A.A. Abu-Hussein, W. Linert, Spectrochim. Acta A 74 (2009) 214–223. [50] A.H. Manikshete, S.K. Sarsamkar, S.A. Deodware, V.N. Kamble, M.R. Asabe, Inorg, Chem. Commun. 14 (2011) 618–621. [51] S. Nagakura, J. Am. Chem. Soc. 80 (1958) 520–524. [52] a K. Watanabe, J. Chem. Phys. 26 (1957) 542–547; b K. Watanabe, J.R. Mottl, ibid. 26 (1957) 1773–1774. [53] W.C. Price, R. Bralsford, P.V. Harris, R.G. Ridley, Spectrochim. Acta 14 (1959) 45–55. [54] A.J.C. Nicholson, J. Chem. Phys. 43 (1965) 1171–1177.