Syntheses, crystal structure and chromotropic properties of dinuclear copper(II) complexes of tertiary diamine with hydroxo bridge

Polyhedron 28 (2009) 3685–3690

Contents lists available at ScienceDirect

Polyhedron journal homepage: www.elsevier.com/locate/poly

Syntheses, crystal structure and chromotropic properties of dinuclear copper(II) complexes of tertiary diamine with hydroxo bridge Hamid Golchoubian *, Roja Zamen Zarabi Dept. of Chem., Univ. of Mazandaran, Babol-sar, Iran

a r t i c l e

i n f o

Article history: Received 10 July 2009 Accepted 4 August 2009 Available online 13 August 2009 Keywords: Tertiary diamine Copper(II) complex Solvatochromism Thermochromism Diamine-ligand Dinuclear complex

a b s t r a c t Dinuclear copper(II) complex with the general formula [(diamine)Cu(OH)]2(ClO4)2, where diamine stands for N,N-dialkyl,N0 -benzyl-ethylenediamine, were prepared. The complexes were characterized by elemental analysis, spectroscopic, conductance measurements, and X-ray structural analysis. The complexes are soluble in various organic solvents and show positive solvatochromism. Thermochromism was also observed particularly in strong donor solvents. Ó 2009 Elsevier Ltd. All rights reserved.

1. Introduction Spectra are a plentiful source of valuable information on the factors influencing the structure and the reactivity of chemical entities. Spectral investigation of copper(II) chelates of ethylenediamine and its derivatives have made a significant contribution towards the understanding of factors determining the stereochemistry, the stability and properties of these important classes of coordination compounds [1–10]. One of the most beautiful properties the copper(II) chelate compounds is chromotropmism in which defined as a reversible variation of the spectral properties of compound under differing physical or chemical conditions such as temperature (thermochromism) and solvent (solvatochromism). There are many applications for this phenomenon such as Lewis acid–base color indicator [11], imaging, [12–15] photo-switching, [16–18] and sensor materials [19]. Several investigations were reported on the formation of secondary [20], tertiary and quaternary diamines [21,22]. These investigations demonstrated that in the case of diamine nitrogen-donor atoms of the ligand with bulky substituents prevent the stabilization of homoleptic bis- or tris-diamine complexes, but the ligand forms a bridging dinuclear l-hydroxo-complex [23]. In continuation of our study on the chromotropic behavior of copper(II) complexes [24,25] we prepared two dinuclear copper(II) complexes abbreviated as [(diamine)Cu(OH)]2(ClO4)2, shown in Scheme 1. * Corresponding author. Tel./fax: +98 1125342350. E-mail address: [email protected] (H. Golchoubian). 0277-5387/$ - see front matter Ó 2009 Elsevier Ltd. All rights reserved. doi:10.1016/j.poly.2009.08.006

These complexes show interesting chromotropism phenomena in solution. We also intend to study effect of steric hindrance around the copper ion imposed by ethylenediamine moiety on the solvatochromic property of the dinuclear copper(II) complexes. Additionally, the X-ray single crystal structure of one of these complexes is discussed in related to its solvatochromic property.

2. Experimental 2.1. Materials and measurements N,N-dimethyl,N0 -benzyl-ethylenediamine and its dinuclear Cu(II) complex were prepared according to published procedure [26]. All solvents were spectral-grade and all other reagents were used as received. All the samples were dried to constant weight under a high vacuum prior to analysis. Caution: perchlorate salts are potentially explosive and should be handled with appropriate care. Conductance measurements were made at 25 °C with a Jenway 400 conductance meter on 1.00  103 M samples in selected solvents. Infrared spectra (potassium bromide disk) were recorded using a Bruker FT-IR instrument. The electronic absorption spectra were measured using a Braic2100 model UV–Vis spectrophotometer. 1H NMR and 13C NMR spectra were recorded on a Bruker 500 DRX Spectrometer. Elemental analyses were performed on a LECO 600 CHN elemental analyzer. Absolute metal percentages were determined by a Varian-spectra A-30/40 atomic absorption-flame spectrometer.

3686

H. Golchoubian, R.Z. Zarabi / Polyhedron 28 (2009) 3685–3690 Table 1 Crystal data and structure refinement for the complex of 1.

R

Empirical formula Formula weight Color Temperature (K) Wavelength (Å) Crystal system Space group Unit cell dimensions a (Å) b (Å) c (Å) b (°) V (Å3) Z Calculated density (g/cm3) Absorption coefficient (mm1) F(0 0 0) Crystal size (mm3) h Range for data collection (°) Index ranges

R H O

N Cu NH

NH Cu

O H

(ClO4)2

N R

R

1: R = CH 3 2: R = C 2H5

l (mm1)

Scheme 1. The structure of dinuclear copper(II) complexes.

Reflections collected/unique [R(int)] Completeness to 2h = 29.00 Refinement method Data/restraints/parameters Final R indicesa [I > 2r(I)]b Goodness-of-fit (GOF) on F2c R indices (all data) Extinction coefficient Largest difference peak and hole (e Å3)

2.2. Syntheses 2.2.1. Preparation of N,N-diethyl,N0 -benzyl-1,2-diaminoethane A typical procedure is as follow: a mixture of benzaldehyde (30 mmol) and N,N-dimethylethylenediamine (30 mmol) in THF (50 mL) was stirred for 24 h. Then the solvent was evaporated under reduced pressure. The resultant yellow oil was then dissolved in n-hexane (40 mL), washed with a minimal volume of water, and dried over anhydrous Na2SO4. Concentration under reduced pressure gave diimine product as yellow oil. To the stirred solution of yellow oil (3.4 g) in methanol (60 mL) at room temperature was gradually added NaBH4 (30 mmol) over 0.5 h. The resulting mixture was allowed to stand overnight, and then was added acetic acid (10 mL). The mixture was made alkaline by NaOH (4 M) and was extracted with dichloromethane (5  10 mL). The combined CH2Cl2 fractions were dried over anhydrous Na2SO4. Evaporation of the solvent under reduced pressure resulted in desired product as orange oil. The typical yields were 60–70%. Selected IR data (m/cm1 using KBr): 3325 (m, N–H str.), 2815 (s, C–H str. aliphatic), 1472 (s, Ph–O str.), 680 (m, C–CH3 + ring def. + Cu–O). 1H NMR (500 MHz, CDCl3), d: 0.91 (t, J = 7.1 Hz, 6H, CH3–CH2–); 2.40 (q, J = 7.1 Hz, 4H, Me–CH2– N–); 2.47 (t, J = 5.4 Hz, 2H, CH2–NH); 2.57 (t, J = 5.4 Hz, 2H, –CH2–N– Et2); 3.70 (s, 2H, CH2 –Ph); 7.12–7.22 (m, 5H, Ar–H). 13C NMR (126.77 MHz in CDCl3) d: 11.8 (CH3–CH2–); 46.7 (Me–CH2–N–); 46.9 (CH2–Ph); 52.5 (CH2–NH); 53.9 (–CH2–N–Et2); 126.7 (CH, pAr); 128.0 (CH, o-Ar); 128.2 (CH, m-Ar); 140.4 (C, Ar). 2.2.2. Preparation dinuclear complexes of copper(II) A typical procedure is as follow: to the solution of desired diamine (10 mmol) in ethanol (30 mL) were slowly added Cu(ClO4)26H2O (5 mmol) in ethanol (10 mL). The resultant mixture was stirred for 2 h at room temperature. The desired compound precipitated from the reaction mixture as violet solid. The typical yield was 80%. Selected IR data (m/cm1 KBr disk): 3530 (m, O–H str.), 3285 (m, N–H str.), 1483 (m, CH2–Ph str.), 1095 (s, Cl–O str.), 695 (w, CH2 rock.), 610 (m, Cl–O bend.), 510 (s, Cu–O str.), 490 (m, Cu-N str.). Anal. Calc. for C26H46N4Cl2O10Cu2: C, 40.42; H, 6.00; N, 7.25; Cu, 16.45. Found: C, 40.31; H, 6.10; N, 7.21; Cu, 16.36%.

a b c

C22H38Cl2Cu2N4O10 716.54 prism, black 100(1) 0.71073 monoclinic P21/c 7.9004(4) 18.3740(10) 10.7504(6) 111.253(5) 1454.41(14) 2 1.363 1.705 740 0.35  0.25  0.15 2.22–29.00 10  h  10 25  k  24 14  l  14 1.71 15 730/3860 [0.0352] 99.8% full-matrix least-squares on F2 3860/0/201 R1 = 0.0299, wR2 = 0.0724 1.000 R1 = 0.0364, wR2 = 0.0757 0.023(3) 0.746 and 0.423

P P ||Fo|  |Fc||/ |Fo|. P 2 2 2 P wR ¼ ½ð ½F o  F c Þ = ½wðF 2o Þ2 1=2 . P S ¼ ½wðF 2o  F 2c Þ2 =ðN obs  N param Þ1=2 . R=

(Bruker) diffractometer with graphite-monochromated Mo Ka radiation (k = 0.71073 Å). Data reduction, including the absorption correction, was performed with the SAINTPLUS (Bruker, 1998a) software package [27]. Solution, refinement and analysis of the structure were performed by using SHELXTL programs [28,29]. The structure was solved by direct methods (SIR92) [30] and refined by the full-matrix least-squares method based on F2 against all reflections [31]. Geometrical calculations were carried out with PLATON [32] and the figures were made by the use of the ORTEP-3 [33] and MERCURY [34] programs. The complete conditions of data collection and structure refinements are given in Table 1. The hydrogen atoms of NH and OH groups were found in difference Fourier synthesis. The H(C) atom positions were calculated. All hydrogen atoms were refined in isotropic approximation in riding model with the Uiso(H) parameters equal to 1.2 Ueq(Ci), for methyl groups equal to 1.5 Ueq(Cii), where U(Ci) and U(Cii) are, respectively, the equivalent thermal parameters of the carbon atoms to which corresponding H atoms are bonded. Refinement of F2 was against all reflections. The weighted R-factor wR and goodness of fit S are based on F2, conventional R-factors R are based on F, with F set to zero for negative F2. The threshold expression of F2 > 2r(F2) is used only for calculating R-factors(gt) etc., and is not relevant to the choice of reflections for refinement. R-factors based on F2 are statistically about twice as large as those based on F, and R-factors based on all data will be even larger.

3. Results and discussion 2.3. X-ray crystallographic study A suitable single crystal of 1 was glued on the tip of a glass fibre. The X-ray data were collected at 100(1) K by x-scans on SMART

The diamine chelate of N,N-diethyl,N0 -benzyl-ethylenediamine was synthesized with similar procedures reported by us [26] except that tetrahydrofuran was found to be better solvent than eth-

H. Golchoubian, R.Z. Zarabi / Polyhedron 28 (2009) 3685–3690

anol so that the yield was improved up to 70%. The dinuclear complexes of copper(II) was prepared with the same procedure reported previously [26]. 3.1. IR spectra The IR spectra of free ligands show strong to medium bands between 2820 and 2760 cm1, which are characteristic of the C–H stretching vibrations of the N–CH3 (or C2H5) group [8], provided that the lone-pair of electrons on the nitrogen atom is not involved in a bond but upon coordination, these bands disappear. Presence of a sharp signal at around 3530 cm1 in both complexes can be assigned to the O–H stretching frequency of the hydroxo bridge [7]. Dependence on coordination is also exhibited by the intense and narrow band occurring at 3250 ± 30 cm1 which is associated with N–H vibration and is observed at around 3280 cm1 and broader in the free diamine ligands. As the lone-pair of electrons of the donor nitrogen atoms become involved in the metal–ligand bond, the transfer of electron density to the metal and the subsequent polarization of the ligands involves electron depopulation of the N–H bond which culminates in a shift to lower frequencies [8]. At lower frequency the complexes also exhibited bands around 503–512 and 480–490 cm1 which are attributed to the m(Cu–O) and

3687

m(Cu–N) vibration modes, respectively [35]. Due to the larger dipole moment change for the Cu–O band compared to the Cu–N band, the m(Cu–O) band usually appears at higher frequency than the m(Cu–N) band [36]. The presence of the ClO4  group is declared by an intense band at around 1090 cm1 and a medium unsplit band at 610 cm1 which are attributed to the anti-symmetric stretching and anti-symmetric bending vibration modes, respectively [37]. The former band at 1090 cm1 is split with a poorly defined maximum showing the deformation from Td symmetry. The infrared spectra of two complexes are very similar with minor change in band maxima except that a extra band is present in the fingerprint region of N,N-diethyl,N0 -benzyl-1,2-diaminoethane. Infrared spectra of the free N,N-diethyl,N0 -benzyl-1,2-diaminoethane ligand and its dinuclear copper(II) complex are illustrated in Fig. 1. 3.2. Conductometric data Molar conductance of the dinuclear complexes measured in some organic solvents with different donor numbers are presented in Table 2. The standard values of 1:2 electrolyte in the respective solvents are shown in the same table [38]. It can be seen that the ionization of ClO4  from the complexes in nitromethane (NM), ace-

Fig. 1. Infrared spectra of free ligand N,N-diethyl,N0 -benzyl-1,2-diaminoethane and its dinuclear Cu(II) complex, 2.

3688

H. Golchoubian, R.Z. Zarabi / Polyhedron 28 (2009) 3685–3690

solvents for coordination and dissociation of ClO4  ions. These data support the chromotropic behavior of the complexes.

Table 2 Molar conductance of the complexes in various solvents.a Complex

1 2 1:2 electrolyte a

KM (ohm1 cm2 mole1) DCM

NM

ACN

AC

MeOH

DMF

2 3 60– 75

85 123 150– 180

117 182 220– 300

140 108 160– 200

127 152 160– 220

155 155 130– 170

Concentration: ca. 1.0  103 M.

Fig. 2. An ORTEP view of compound 1 in representation of atoms via thermal ellipsoids at 50% probability level. Atoms labeled with A are obtained from the basic ones by the symmetry operation x, y + 1, z.

tone (AC), acetonitrile (ACN) or methanol (MeOH) is not complete, which indicates partial dissociation into ClO4  anion and the cationic planar complexes. In dichloromethane (DCM), the perchlorate complexes demonstrate almost non-electrolyte behavior, which means that the species in solution are the same as that of the solid state. However, in a very strong donor solvent like dimethylformamide (DMF), the conductivity data show that the perchlorate complexes almost consistent with 2:1 electrolytes. These results are similar with those were observed before for mixed-chelate copper(II) complexes [25,26]. Consequently, the complexes are strongly influenced by the donor property of the

3.3. X-ray structure The X-ray analysis of compound 1 was determined. An ORTEP view of this compound is shown in Fig. 2 together with the numbering scheme. The selected bond lengths and angles are given in Table 3. The complex is a symmetric dinuclear unit made up of two distorted coordination square pyramids. The two diamine nitrogen atoms and two oxygen atoms of hydroxo bridges occupying the basal plane and the oxygen atom of the perchlorate group in the apical position with significantly longer bond length of 2.5219(15) Å. According to Addison and Rao [39] the distortion of the square pyramidal geometry towards trigonal bipyramidal can be described by geometrical parameter s = |b  a|/60, where b and a are the bond angles involving the trans donor atoms in the basal plane. The s value for the coordination around the copper atom is 0.053, confirming the square pyramidal geometry. The basal atoms are nearly coplanar; the deviations from the least-squares plane through the CuN2O2 atoms are N(1) 0.035, N(2) 0.032, O(1W) 0.037, O(1WA) 0.036, Cu(1) 0.005 Å. The mean Cu– N(amine) distance of 2.03 Å and the bite angle N(1)–Cu(1)–N(2) of 85.29(6)° are close to the corresponding average values of the copper(II) complexes with ethylenediamine reported [40,41]. The Cu(1)–O(1W) and Cu(1)–O(1WA) distances are 1.926(1) Å and 1.949(1) Å, respectively. The shortest interdinuclear Cu(1)  Cu(1A) separation is 2.945(3) Å. The phenyl groups are directed away from the N2CuO2CuN2 plane and in opposite side of perchlorate groups to minimize the steric hindrance around the copper(II) ions. The five-membered chelate ring in the complex is puckered so that the torsion angle of N(1)–C(1)–C(2)–N(2) is 45.34(18)°. The perchlorate ions positioned in apical site of the basal plane are loosely coordinated to copper(II) ions and can be driven out from the coordination sphere by solvent molecules in solution, which will be more sensible bonded with the increase in their donor power, leading to the observed solvatochromism. Donor nitrogen atoms N(1) of amine group, and oxygen atoms O(1) hydroxo group form hydrogen bonds to adjacent acceptor oxygen atoms O(1) and O(2) of perchlorate anions. In the crystal lattice, the dinuclear units are packed through intermolecular N–H  O and O–H  O hydrogen bonding interactions with the ClO4  anions. The hydrogen bonds form macrocycle-like ring systems from the interactions of N(1)–H and O(1W)–H with O(1) atom of the neighbor molecules (dotted lines in Fig 3).

Table 3 Selected bond lengths (Å) and angles (°) for complex 1. Bond distances 2.0395(14) 2.5219(15) 2.945(3)

Cu(1)–N(2) Cu(1)–O(4) Cu(1)–Cu(1A)

1.9262(12) 1.9489(12) 2.0327(14)

Cu(1)–O(1WA) Cu(1)–O(1W) Cu(1)–N(1)

Bond angles 85.29(6) 87.27(5) 85.59(5) 89.96(6) 96.07(6)

N(1)–Cu(1)–N(2) O(1WA)–Cu(1)–O(4) O(1W)–Cu(1)–O(4) N(1)–Cu(1)–O(4) N(2)–Cu(1)–O(4)

81.06(5) 175.15(5) 94.75(5) 98.97(5) 178.34(5)

O(1WA)–Cu(1)–O(1W) O(1WA)–Cu(1)–N(1) O(1W)–Cu(1)–N(1) O(1WA)–Cu(1)–N(2) O(1W)–Cu(1)–N(2)

Torsion angles 18.86(17) 45.34(18) 72.47(17) 0.0

Cu(1)–N(1)–C(1)–C(2) N(1)–C(1)–C(2)–N(2) C(10)–N(2)–C(2)–C(1) O(1WA)–Cu(1)–O(1W)–Cu(1A)

177.52(6) 47.90(15) 70.95(17) 57.09(16)

N(1)–Cu(1)–O(1W)–Cu(1A) Cu(1)–N(2)–C(2)–C(1) C(1)–N(1)–C(3)–C(4) Cu(1)–N(1)–C(3)–C(4)

3689

H. Golchoubian, R.Z. Zarabi / Polyhedron 28 (2009) 3685–3690

1.2 1

Abs

0.8 0.6

Py DCM CH3CN MeOH

0.4 0.2

Fig. 3. The intermolecular hydrogen bonding in the dimeric structure of 1.

DMF

0

3.4. Solvatochromism

12

16

ν /10

3

18

20

-1

cm

Fig. 4. Absorption spectra of the complex 1 (5.0  103 M) in selected solvents. Absorption spectra in other solvents are omitted for clarity.

18

νmax (x103 cm-1)

The complexes are soluble in a large number of organic solvents, forming violet solutions in solvents of low donor number (DN) [42] and changing through blue to green as the solvent DN increases. The observed mmax values of d–d bands of the complexes in various solvents are summarized in Table 4 and plotted versus the DN of the respective solvents in Fig 4. The mmax values decreases almost linearly with the increasing DNs of solvents. This originates in variation of Lewis acid–base interaction between the chelate ion and the respective solvent molecules. Since all the dxy, d yz and dxz orbitals of the Cu(II) ion are raised up by its interaction with polar solvent molecules approaching from above and below of the molecular plane, the broad d–d transition band of complexes moves to the red with the increase of the DN of the solvent. Regression analysis of band maxima of compounds 1 and 2 against donor number of solvents is shown in Fig. 5 and indicated good correlation and also confirms the solvatochromic behavior of these complexes. Deviations were observed indicative of the complex nature of the solute-solvent interaction. For instants, the kmax of these complexes in CH2Cl2 occur at longer wavelength than in nitromethane, although the donor number of CH2Cl2 (= 0) is lower than that of nitromethane (= 2.7). This anomaly was ascribed to the formation of ion pairs by mean of an axial coordination of ClO4  in CH2Cl2. As the relative dielectric constant of nitromethane takes a much higher value (28.5 for nitromethane versus 8.9 for CH2Cl2 in room temperature), this solvent facilitates the dissociation of cationic chelate. Thus reducing the axial interactions of the ClO4  and eliminating the tetragonal distortion [43].

14

1: R 2 = 0.97 2: R 2 = 0.93 17 ♦ 1 ■ 2

16

15 0

5

10

15

20

25

30

35

DN Fig. 5. Dependence of the mmax values of compounds 1 and 2 on the solvent donor number values.

1.2 1

3.5. Thermochromism In solution, both complexes demonstrate irreversible thermochromism in strong donor solvents. The original blue color solution gradually losses its intensity and turn yellow on heating as illustrated in Fig. 6. This seems to be due to the decomposition of the chelate cation caused by solvent attack. In case of pyridine, it seems that the diamine ligands are driven off, since the addition of free ligand to the solution apparently suppresses the decomposition so that the yellow color appears in longer time. More over, characterization of the product obtained after heating in the case of pyridine as a solvent demonstrated the lack of diamine ligand. Although the exact mechanism of this thermochromism is still un-

Abs

0.8

a

0.6 b

0.4 c d

0.2 0 12

14

16

18

ν /103 cm-1

20

22

Fig. 6. Spectral changes of complex 1 (5.0  103 M) in solvent of pyridine at 50 °C after (a) 0 min, (b) 10 min, (c) 20 min and (d) 30 min.

Table 4 Electronic spectra of the complexes in various solvents. Solvent DN

mmax/103 cm1 (e/M1 cm1)

Complex 1 Complex 2

DCM 0.0

MeNO2 2.7

ACN 14.1

MeOH 23.5

DMF 26.6

Py 33.1

17.71 (172) 17.69 (145)

17.74 (149) 17.66 (137)

16.88 (145) 16.89 (126)

16.86 (180) 16.94 (196)

16.04 (165) 16.26 (120)

15.54 (150) 15.14 (115)

3690

H. Golchoubian, R.Z. Zarabi / Polyhedron 28 (2009) 3685–3690

known these facts may indicate that equilibrium showed in (1) is shifted to the right-hand side by heating. þ ½ðdiamineÞCuðOHÞ2þ 2 þ nPy 2½CuðOHÞðPyÞn  þ 2 diamine

ð1Þ

4. Conclusion Two dinuclear copper(II) complexes have been synthesized and characterized in combination with two new organic ligands. X-ray structure determination demonstrates square pyramidal geometry around the metal centers. The crystal structure is stabilized by N–H  O and O–H  O intermolecular hydrogen bonding of the oxygen atoms of perchlorate ion with the nitrogen atom of the secondary amine and the oxygen atom of the hydroxo bridges and Van der Waals interaction. However, in solution, both complexes exhibit solvatochromism, which is due to substitution of perchlorate groups by the solvent molecules. The change of the geometry as well as ligand field strength around the metal center leads to change of absorption band of the complexes and the shift of the d–d bands in different solvent media. The solvatochromic properties of two dinuclear copper(II) complexes are almost similar. As a result, the steric hindrance around the copper ion does not play substantial role in the solvatochromic properties of these complexes. They also demonstrate theromochromism behavior in strong donor power solvents. Acknowledgement We are grateful for the financial support of Mazandaran University of the Islamic Republic of Iran. Appendix A. Supplementary material CCDC 724016 contains the supplementary crystallographic data for this paper. These data can be obtained free of charge via http:// www.ccdc.cam.ac.uk/conts/retrieving.html, or from the Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: (+44) 1223-336-033; or e-mail: [email protected]. Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.poly.2009.08.006. References [1] Y.-B. Jiang, H.-Z. Kou, R.-J. Wang, A.-L. Cui, Eur. J. Inorg. Chem. (2004) 4608. [2] U. El-Ayann, F. Murata, Y. Fukuda, Monatsh. Chem. 132 (2001) 1279.

[3] H. Miyamae, H. Kudo, G. Hihara, K. Sone, Bull. Chem. Soc. Jpn. 71 (1998) 2621. [4] S. Haddad, R.D. Willett, Inorg. Chem. 40 (2001) 2457. [5] H.-D. Bian, J.-Y. Xu, W. Gu, S.-P. Yan, P. Cheng, D.-Z. Liao, Z.-H. Jiang, Polyhedron 22 (2003) 2927. [6] S.F. Pavkovic, D.E. Meek, Inorg. Chem. 4 (1965) 20. [7] D.E. Meek, S.A. Ehrhardt, Inorg. Chem. 4 (1965) 584. [8] C. Tsiamis, M. Themeli, Inorg. Chim. Acta 206 (1993) 105. [9] Y. Fukada, M. Cho, K. Sone, Bull. Chem. Soc. Jpn. 62 (1989) 51. [10] K.C. Patel, R.R. Patel, J. Inorg. Nucl. Chem. 39 (1977) 1325. [11] K. Miyamoto, M. Sakamoto, C. Tanaka, E. Horn, Y. Fukuda, Bull. Chem. Soc. Jpn. 78 (2005) 1061. [12] R. Yamamoto, T. Maruyama, Jpn. Kokai Tokyo Koho JP 06073169, 1994. [13] Y. Masuda, F. Inatsugu, Jpn. Kokai Tokyo Koho JP 50121176, 1975. [14] M.L. Golden, J.C. Yarbrough, J.H. Reibenspies, N. Bhuvanesh, P.L. Lee, Y. Zhang, M.Y. Darensbourg, in: 228th ACS National Meeting, Philadelphia, PA, United State, August 22–26, 2004. [15] C.A. Fontan, R.A. Olsina, Talanta 36 (1989) 945. [16] O. Sato, Acc. Chem. Rev. 36 (2003) 692. [17] S.R. Marder, J.W. Perry, Science 263 (1994) 1706. [18] O. Sato, S. Hayami, Y. Einaga, Z.-Z. Gu, Bull. Chem. Soc. Jpn. 76 (2003) 443. [19] K.S. Schanze, R.H. Schmehl, J. Chem. Educ. 74 (1997) 633. [20] L.M. Mirica, D.J. Rudd, M.A. Vance, E.I. Solomon, K.O. Hodgson, B. Hedman, T. D.P. Stack, J. Am. Chem. Soc. 128 (2006) 2654. [21] P. Akilan, M. Thirumavalavan, M. Kandaswamy, Polyhedron 22 (2003) 3483. [22] M. Thirumavalavan, P. Akilan, M. Kandaswamy, Inorg. Chem. Commun. 5 (2002) 422. [23] F. Murata, M. Arakawa, A. Nakao, K. Satoh, Y. Fukuda, Polyhedron 26 (2007) 1570. [24] E. Movahedi, H. Golchoubian, J. Mol. Struct. 787 (2006) 167. [25] (a) H. Asadi, H. Golchoubian, R. Welter, J. Mol. Struct. 779 (2005) 30; (b) H. Golchoubian, E. Rezaee, J. Mol. Struct. 927 (2009) 91. [26] H. Golchoubian, E. Rezaee, J. Mol. Struct. 929 (2009) 154. [27] Bruker, SAINTPLUS, Data Reduction and Correction Program v. 6.01, Bruker AXS, Madison, Wisconsin, USA, 1998. [28] G.M. Sheldrick, SADABS v.2.01, Bruker/Siemens Area Detector Absorption Correction Program, Bruker AXS, Madison, Wisconsin, USA, 1998. [29] G.M. Sheldrick, SHELXTL v. 5.10, Structure Determination Software Suite, Bruker AXS, Madison, Wisconsin, USA, 1998. [30] A. Altomare, G. Cascarano, C. Giacovazzo, A. Guagliardi, J. Appl. Crystallogr. 26 (1993) 343. [31] G.M. Sheldrick, SHELXL-98b, Program for the Refinement of Crystal Structures, Bruker AXS, Madison, Wisconsin, USA, 1998. [32] A.L. Spek, J. Appl. Crystallogr. 36 (2003) 7. [33] L.J. Farrugia, J. Appl. Crystallogr. 30 (1997) 568. [34] C.F. Macrae, P.R. Edgington, P. McCabe, E. Pidcock, G.P. Shields, R. Taylor, M. Towler, J. Van de Streek, J. Appl. Crystallogr. 39 (2006) 453. [35] N. Raman, S. Esthar, C. Thangaraja, J. Chem. Sci. 116 (2004) 209. [36] N. Chkaku, K. Nakamoto, Inorg. Chem. 10 (1971) 798. [37] Y. Fukuda, A. Shimura, M. Mukaida, E. Fujita, K. Sone, J. Inorg. Nucl. Chem. 36 (1974) 1265. [38] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81. [39] A.W. Addison, T.N. Rao, J. Chem. Soc., Dalton Trans. (1984) 1349. [40] H. Golchoubian, Anal. Sci. 24 (2008) x169. [41] H. Golchoubian, Anal. Sci. 24 (2008) x195. [42] V. Gutman, The Donor–Acceptor Approach to Molecular Interactions, Plenum Press, New York, 1978. [43] R.W. Soukup, K. Sone, Bull. Chem. Soc. Jpn. 60 (1987) 2286.