novel vic-dioxime ligands and their poly-metal complexes bearing 1,8-diamino-3,6-dioxaoctane:...

Novel vic-dioxime ligands and their poly-metal complexesbearing 1,8-diamino-3,6-dioxaoctane: synthesis, characterization,spectroscopy and electrochemistry

Ahmet Kilic Æ Mustafa Durgun Æ Esref Tas ÆIsmail Yilmaz

Received: 12 July 2007 / Accepted: 8 August 2007 / Published online: 22 November 2007

� Springer Science+Business Media B.V. 2007

Abstract Three novel vic-dioxime ligands containing the

1,8-diamino-3,6-dioxaoctane group, N,N0-(1,8-diamino-3,6-

dioxaoctane)-p-tolylglyoxime (L1SL1H4), N,N0-(1,8-dia-

mino-3,6-dioxaoctane)-phenylglyoxime (L2SL2H4), and

N,N0-(1,8-diamino-3,6-dioxaoctane)-glyoxime (L3SL3H4)

have been prepared from 1,8-diamino-3,6-dioxaoctane with

anti-p-tolylchloroglyoxime, anti-phenylchloroglyoxime or

anti-monochloroglyoxime. Polynuclear complexes

[M(LxSLx)]n or [M(LxSLx)(H2O)]n (x = 1, 2 and 3), where

M = CuII, CoII, and NiII, have been obtained with 1:1 metal/

ligand ratio. The CuII and NiII poly-metal complexes of these

ligands are proposed to be square planar, while also the

prepared CoII complexes are proposed to be octahedral with

two water molecules as axial ligands. The detection of H-

bonding in the [Ni(L1SL1)]n, [Ni(L2SL2)]n and

[M(L3SL3)(H2O)]n metal complexes by FT i.r. spectra

revealed the square planar or octahedral [MN4�H2O)]n

coordination of poly-nuclear metal complexes. [MN4]n

coordination of the [Ni(L1SL1)]n and [Ni(L2SL2)]n com-

plexes were also determined by 1H-n.m.r. spectroscopy. The

ligands and poly-metal complexes were characterized by

elemental analyses, FT-i.r., u.v.-vis., 1H and 13C-n.m.r.

spectra, magnetic susceptibility measurements, molar con-

ductivity, cyclic voltammetry, and differential pulse

voltammetric (DPV) techniques.

Introduction

The synthesis of vic-dioximes and their different deriva-

tives have been the subject of study for very a long period

of time. Macromolecules attached to dioximes and their

transition metal complexes have been investigated [1, 2].

The transition metal complexes of vic-dioximes have been

of particular interest as biological model compounds [3].

Recently, there has been growing interest in polymeric

compounds. Polymeric metal complexes of vic-dioximes

show interesting and important characteristics, especially

in areas such as semiconductors [4], heat resistance mate-

rials [5], and gas separators [6]. Moreover, vic-dioximes

are interesting for many applications in a variety of high

technology fields, such as medicine [7–9], catalysis [10,

11], electrooptical sensors [12], liquid crystals [13], and

trace metal analysis [14]. Interest in the metal coordination

environment has prompted the study of oxime ligands due

to their variable geometries [3] and the tunability of their

substituents [15, 16]. The exceptional stability and unique

electronic properties of these complexes can be attributed

to their planar structure, which is stabilized by hydrogen

bonding [17]. The investigations of the redox properties of

these types of complexes are also of great interest in terms

of their various technological applications [18]. The oxi-

dation states of the central metals, type and number of

donor atoms and core structures of the complexes are major

factors in determining structure–function relations of the

transitions metal complexes [11]. Also, the two hydrogen

bridges have been substituted with metal complexes to

obtain polynuclear compounds in order to investigate the

magnetic interactions [19].

The aim of the present study was to synthesize and

characterize three novel vic-dioximes (Scheme 1) contain-

ing 1,8-Diamino-3,6-Dioxaoctane and to obtain their

A. Kilic (&) � M. Durgun � E. Tas

Department of Chemistry, University of Harran, 63190

Sanliurfa, Turkey

e-mail: [email protected]

I. Yilmaz

Department of Chemistry, Technical University of Istanbul,

34469 Istanbul, Turkey

123

Transition Met Chem (2008) 33:29–37

DOI 10.1007/s11243-007-9010-6

polymeric complexes with Co(II), Cu(II), and Ni(II) ions

(Fig. 1). The cyclic voltammetric measurements of the

polymeric complexes have also been studied to understand

electrochemical behavior of their reduced or oxidized spe-

cies in nonaqueous solution. The ligands and poly-metal

complexes were characterized by elemental analyses, FT-

i.r., u.v.-vis., 1H and 13C-n.m.r. spectra, magnetic suscep-

tibility measurements, molar conductivity, cyclic

voltammetry, and differential pulse voltammetric (DPV)

techniques.

Experimental

Material and methods

All reagents and solvents were of reagent-grade quality and

obtained from commercial suppliers. (Fluka Chemical Com-

pany, Taufkirchen, Germany) and tetra-n-butylammonium

perchlorate (TBAP, Fluka Chemical Company, Taufkirchen,

Germany) were used as received. Anti-phenylchloroglyoxime

and anti-p-tolylchloroglyoxime were synthesized as descri-

bed in the literature [20], anti-monochloroglyoxime was

synthesized as described in the literature [21]. The elemental

analyses were carried out in the Laboratory of the Scientific

and Technical Research Council of Turkey (TUBITAK). FT-

i.r. spectra were recorded on a Perkin Elmer Spectrum RXI

FT-i.r. Spectrometer as KBr pellets. 1H-n.m.r. and 13C-n.m.r.

spectra were recorded on a Bruker-Avence 400 MHz spec-

trometer. Magnetic susceptibilities were determined on a

Sherwood Scientific Magnetic Susceptibility Balance (Model

MK1) at room temperature (20 �C) using Hg[Co(SCN)4] as a

calibrant; diamagnetic corrections were calculated from Pas-

cal’s constants [22]. UV-vis spectra were recorded on a

Shimadzu 1601 PC. Molar conductivities (KM) were recorded

on a Inolab Terminal 740 WTW Series. Cyclic voltammo-

grams (c.v.) were carried out using c.v. measurements with

Princeton Applied Research Model 2263 potentiostat con-

trolled by an external PC. A three-electrode system (BAS

model solid cell stand) was used for c.v. measurements in

DMSO and consisted of a 2 mm sized platinum disc electrode

as working electrode, a platinum wire counter electrode, and

an Ag/AgCl reference electrode. The reference electrode was

separated from the bulk solution by a fritted-glass bridge filled

with the solvent/supporting electrolyte mixture. The ferro-

cene/ferrocenium couple (Fc/Fc+) was used as an internal

standard but all potentials in the article are referenced to the

Ag/AgCl reference electrode. Solutions containing com-

plexes were deoxygenated by a stream of high purity nitrogen

for at least 5 min. Before running, the experiment and the

solution was protected from air by a blanket of nitrogen during

the experiment. Controlled potential electrolysis (CPE) was

performed with Princeton Applied Research Model 2263

potentiostat /Galvanostat. An BAS model electrolysis cell

with a fritted-glass to separate the cathodic and anodic por-

tions of the cell was used for bulk electrolysis. The sample and

solvent were placed into the electrolysis cell under nitrogen.

Synthesis of the ligands L1SL1H4, L2SL2H4, and

L3SL3H4

1,8-Diamino-3,6-Dioxaoctane (1.00 g, 6.80 mmol for

L1SL1H4, 1.00 g, 6.80 mmol for L2SL2H4 and 1.00 g,

6.80 mmol for L3SL3H4, respectively) was dissolved in

65 cm3 absolute THF. Then, triethylamine (Et3N) (1.38 g,

13.6 mmol for L1SL1H4, 1.38 g, 13.6 mmol for L2SL2H4, and

1.38 g, 13.6 mmol for L3SL3H4, respectively) was added and

the mixture was cooled to -15 �C and kept at this temperature.

To this solution, anti-p-tolylglyoxime (2.88 g, 13.6 mmol),

anti-phenylchloroglyoxime (2.70 g, 13.6 mmol) or anti-

monochloroglyoxime (1.67 g, 13.6 mmol) in 40 cm3 abso-

lute THF was added dropwise under a N2 atmosphere with

NH

OO

HNC

CN

CNR

C N

N

R

O

O

O

O

H

H

M

n

R: CH3, or H

M: NiII, NiII.2H2O, CuII or CoII.2H2O

Fig. 1 The structure of the NiII, CuII, and CoII metal complexes

NH

OO

HN

C N

C N

R OH

OHCN

CN

RHO

HO

H2NO

ONH2

C=N

C=NOH

OHR

CI

R: CH3, and H

FH

T

-15oC

,(C2 H

5 )3 N

a b

cd

e

cd

f

Scheme 1 Synthesis of ligands (L1SL1H4, L2SL2H4, and L3SL3H4)

30 Transition Met Chem (2008) 33:29–37

123

continuous stirring. The addition of the anti-p-tolylglyoxime,

anti-phenylchloroglyoxime or anti-monochloroglyoxime

solutions were carried out during 4 h. The mixture was stirred

for more than 2 h and the temperature was raised to 25 �C.

Precipitated Et3NHCl was filtered off and the filtrate was

evaporated to remove THF. The oily products were dissolved

in CH2Cl2 (15 cm3) and n-hexane (100 cm3) added to pre-

cipitate the compounds. This process was then repeated

several times. Ligands (L1SL1H4, L2SL2H4 and L3SL3H4)

were filtered off and dried in a vacuum at 35 �C. The ligands

are soluble in common organic solvents such as THF, EtOH,

and DMSO. 1H-n.m.r. of L1SL1H4 (DMSO-d6, TMS, d ppm)

d: 5.72–5.80 (d, 1H, NH, Ja = 3 Hz), 7.53–7.51 (m, 4H, Ar-

CH), 7.22–7.20 (m, 4H, Ar-CH), 3.53–3.44 (m, 4H, O-CH2),

3.38–3.30 (m, 4H, O-CH2), 2.40–2.26 (m, 4H, N-CH2), 1.41–

1.15 (m, 6H, C-CH3). 1H-n.m.r. of L2SL2H4 (DMSO-d6,

TMS, d ppm) d: 11.66 (s, 1H, OH) and 9.82 (s, 1H, OH), 5.70

(s, 2H, NH), 7.56–7.41 (m, 4H, Ar-CH), 7.40–7.36 (m, 6H,

Ar-CH), 3.40–3.31 (m, 4H, O-CH2), 3.25–3.12 (m, 4H, O-

CH2), 2.58–2.43 (m, 4H, N-CH2), and 1H-n.m.r. of L3SL3H4

(DMSO-d6, TMS, d ppm) d: 11.41 (s, 1H, OH) and 10.25 (s,

1H, OH), 7.32 (s, 2H, N=CH), 5.54–5.51 (t, 2H, NH, Ja = 9

Hz), 3.54–3.50 (m, 4H, O-CH2), 3.45–3.43 (m, 4H, O-CH2),

2.51–2.49 (t, 4H, N-CH2, Jb = 6 Hz). 13C-n.m.r. of L1SL1H4

(DMSO-d6, TMS, d ppm): C1,20(21.41), C2,19(129.64),

C3,18(128.88), C4,17(129.23), C5,16(139.06), C6,15(150.92),

C7,14(148.36), C8,13(42.63), C9,12(70.73) and C10,11(69.91).13C-n.m.r. of L2SL2H4 (DMSO-d6, TMS, d ppm):

C1,18(129.56), C2,17(128.27), C3,16(129.28), C4,15(132.21),

C5,14(150.83), C6,13(148.80), C7,12(42.76), C8,11(70.79), and

C9,10(69.93). 13C-n.m.r. of L3SL3H4 (DMSO-d6, TMS, dppm): C1,10(148.85), C2,9(144.39), C3,8(42.91), C4,7(71.03),

and C5,6(70.04).

Synthesis of the NiII, CuII and CoII poly-metal

complexes

A solution of nickel(II) chloride hexahydrate (0.12 g,

0.5 mmol), cobalt(II) chloride hexahydrate (0.12 g,

0.5 mmol), or copper(II) chloride dihydrate (0.085 g,

0.5 mmol) in ethanol (20 cm3) was added to a solution of

L1SL1H4 (0.5 g, 1.0 mmol), L2SL2H4 (0.47 g, 1.0 mmol),

or L3SL3H4 (0.32 g, 1.0 mmol) in ethanol (80 cm3) at 60–

65 �C. A distinct change was observed in color from col-

orless to red, brown or green under a N2 atmosphere with

continuous stirring. Then, a decrease in the pH of the

solution was observed. The pH of the solution was ca. 1.5–

3.0 and was adjusted to 4.5–5.5 by the addition of a 1%

triethylamine solution in EtOH. After heating the mixture

for 2 h in a water bath, the precipitate was filtered off,

washed with H2O and diethyl ether several times, and dried

in vacuo at 35 �C. 1H-n.m.r. of [Ni(L1SL1)]n (CDCl3, TMS,

d ppm) d: 15.53 (s, 2H, O–H��O), 5.80–5.77 (t, 2H, NH,

Ja = 10.5 Hz), 7.38–7.36 (d, 4H, Ar-CH, Jc = 7.5 Hz),

7.20–7.17 (d, 4H, Ar-CH, Jd = 8.1 Hz), 3.58–3.29 (m, 8H,

O-CH2), 2.97–2.91 (m, 2H, N-CH2), 2.65–2.58 (m, 2H, N-

CH2), and 1.31–1.28 (m, 6H, C-CH3). 1H-n.m.r. of

[Ni(L2SL2)]n (CDCl3, TMS, d ppm) d: 15.43 (s, 2H, O–

H��O), 5.81–5.78 (d, 2H, NH, Je = 4.8 Hz), 7.57–7.49 (s,

4H, Ar-CH), 7.39–7.38 (d, 6H, Ar-CH, Jc = 2.7 Hz), 3.58–

3.51 (t, 4H, O-CH2, Jf = 11.7 Hz), 3.43–3.29 (m, 4H, O-

CH2), 2.93–2.89 (m, 2H, N-CH2) and 2.59–2.56 (m, 2H, N-

CH2).

Results and discussion

Synthesis

The synthetic route for the ligands (L1SL1H4, L2SL2H4 and

L3SL3H4) is given in Scheme 1. The analytical data of

[M(LxSLx)]n or [M(LxSLx)(H2O)]n (x = 1, 2 and 3) poly-

metal complexes indicate 1:1 metal/ligand stoichiometry.

Additional analytical data are given in Tables 1–4 and the

experimental section.

N.m.r. spectra

The 1H-n.m.r. spectra of the ligands (L1SL1H4, L2SL2H4,

and L3SL3H4) in DMSO-d6 and the Ni(II) complexes in

CHCl3-d1 are given in the experimental section. Although,

no chemical shifts were observed for =N-OH groups of

oximes in the 1H-n.m.r. spectra of L1SL1H4, the other

signals (L1SL1H4) support the structures (in experimental

section). The proton resonance appears as two peaks, low-

intensity singlets at 11.66 and 9.82 ppm for L2SL2H4, at

11.41 and 10.25 ppm for L3SL3H4, respectively. These two

D2O-exchangeable singlets correspond to two non-equiv-

alent -OH protons, which also indicate the anti-

configuration of the -OH relative to each other [23, 24].

The first one is assigned to the -OH proton on the phenyl

side with the latter to the -OH proton of the amidoxime

group, since the effect of various substituents is expected to

be higher on the amidoxime group [25]. The chemical

shifts which belong to -NH protons were observed at 5.71–

5.70 ppm as a doublet for L1SL1H4 at 5.70 ppm as a singlet

for L2SL2H4 and at 5.54–5.51 ppm as a triplet for

L3SL3H4, respectively and disappeared with D2O

exchange. Also, the chemical shifts which belong to

CH=N- protons were observed at 7.32 ppm as a singlet for

L1SL1H4. The 1H-n.m.r. spectra of the diamagnetic metal

complexes, [Ni(L1SL1)]n and [Ni(L2SL2)]n were charac-

terized by the absence of the =N-OH signals in the ligands

and the existence of intra-molecular D2O-exchangeable

Transition Met Chem (2008) 33:29–37 31

123

H-bridge protons was observed by new signals at low field,

d = 15.53 ppm for [Ni(L1SL1)]n and d = 15.43 ppm for

[Ni(L2SL2)]n. Consequently, we may conclude that both of

these d8 metal ions are coordinated with the dioximato

donor sites in square planar geometry [14, 26]. The carbon

resonances of oxime groups are found at 150.92 and

148.36 ppm for L1SL1H4, 150.83 and 148.80 ppm for

L2SL2H4, and 148.85 and 144.39 ppm for L3SL3H4,

respectively. These non-equivalent carbon atoms, particu-

larly, belong to hydroxyimino carbon atoms, also confirm

the anti-structure of L1SL1H4, L2SL2H4, and L3SL3H4,

respectively [19c, 27].

The elemental analysis result shows that the polymer

complexes have been obtained with 1:1 metal/ligand ratio.

Best evidence of the polymer formation is the disappear-

ance of the =N-OH signals in the ligands and the existence

of intra-molecular D2O-exchangeable H-bridge (O–H��O)

protons, which were observed by a new signals at

d = 15.53 ppm for [Ni(L1SL1)]n and d = 15.43 ppm for

[Ni(L2SL2)]n. If the complexation had resulted in the oli-

gomer formation, the 1H NMR spectra of the oligomer

would show both the signals of the H-bridge (O–H��O) at

low field and free C=N-OH protons of oximes due to the

existence of the free oximes at the ends of the oligomer.

The free oximes in the oligomer is expected to appear in

the 1H NMR spectra compared with those of the polymer

complexes, but the corresponding free C=N-OH protons

was not observed in the polymer complexes. Since, the

polymer complexes have high molecular weight, the signal

of the free oximes at the ends of the polymer is probably

not detected.

FT- i.r. Spectra

The infrared spectra of ligands (L1SL1H4, L2SL2H4, and

L3SL3H4) with their Ni(II), Cu(II), and Co(II) metal

complexes have been studied in order to characterize their

structures. The i.r. spectra of Ni(II), Cu(II), and Co(II)

poly-metal complexes were interpreted by comparing the

spectra with those of the free ligands. The characteristic

infrared data are given in Table 2. In the i.r. spectra of the

free ligands, the t(O–H/ N–H) stretching vibrations were

observed at between 3605 and 3060 cm-1 for L1SL1H4,

3629–3084 cm-1 for L2SL2H4 and 3546–3149 for

L3SL3H4, respectively. The free vic-dioxime ligands

showed a strong peak at 1640 cm-1 for L1SL1H4, at

1641 cm-1 for L2SL2H4, and at 1647 cm-1 for L3SL3H4

which is characteristic of the azomethine t(C=N) group

[28]. The t(C=N) stretching vibrations are affected upon

complexation and are situated at a frequency significantly

different than the free ligands. Coordination of the vic-

dioxime ligands to the metal center through the fourTa

ble

1T

he

form

ula

,fo

rmu

law

eig

ht,

colo

rs,

mel

tin

gp

oin

ts,

mo

lar

con

du

ctiv

ity

,y

ield

s,m

agn

etic

susc

epti

bil

itie

s,an

del

emen

tal

anal

yse

sre

sult

so

fth

eco

mp

ou

nd

s

Co

mp

ou

nd

sF

.W(g

/mo

l)C

olo

rD

eco

m.

po

int

(�C

)K

M(X

-1cm

2m

ol-

1)

Yie

ld(%

)l

eff

(BM

)E

lem

ente

lan

aly

ses

%ca

lcu

late

d(f

ou

nd

)

CH

N

Lig

and

(L1S

L1H

4)

C24H

32N

6O

65

00

Yel

low

94

-8

6–

57

.6(5

7.3

)6

.4(6

.3)

16

.8(1

6.6

)

[Co

(L1S

L1)(

H2O

)]n

C24H

34N

6O

8C

o5

93

Bro

wn

10

82

36

33

.96

48

.6(4

8.3

)5

.7(5

.5)

14

.2(1

4.0

)

[Cu

(L1S

L1)]

nC

24H

30N

6O

6C

u5

62

Dar

kg

reen

14

93

66

41

.84

51

.2(5

1.1

)5

.3(5

.2)

15

.0(1

4.8

)

[Ni(

L1S

L1)]

nC

24H

30N

6O

6N

i5

57

Red

24

52

87

0D

ia5

1.7

(51

.4)

5.4

(5.2

)1

5.1

(15

.0)

Lig

and

(L2S

L2H

4)

C22H

28N

6O

64

72

Yel

low

11

4–

82

–5

5.9

(55

.6)

6.0

(5.7

)1

7.8

(17

.3)

[Co

(L2S

L2)(

H2O

)]n

C22H

30N

6O

8C

o5

65

Bro

wn

13

63

66

13

.82

46

.7(4

6.6

)5

.3(5

.2)

14

.9(1

4.9

)

[Cu

(L2S

L2)]

nC

22H

26N

6O

6C

u5

34

Dar

kg

reen

14

25

36

61

.79

49

.4(4

9.3

)4

.9(4

.7)

15

.7(1

5.8

)

[Ni(

L2S

L2)]

nC

22H

26N

6O

6N

i5

29

Red

21

23

47

2D

ia4

9.9

(49

.8)

4.9

(4.9

)1

5.9

(16

.0)

Lig

and

(L3S

L3H

4)

C10H

20N

6O

63

20

Yel

low

14

9–

76

–3

7.5

(37

.2)

6.3

(6.1

)2

6.2

(26

.0)

[Co

(L3S

L3)(

H2O

)]n

C10H

22N

6O

8C

o4

13

Bro

wn

17

83

45

83

.86

29

.1(2

8.9

)5

.3(5

.2)

20

.3(2

0.2

)

[Cu

(L3S

L3)]

nC

10H

18N

6O

6C

u3

82

Dar

kg

reen

16

84

16

11

.81

31

.4(3

1.3

)4

.7(4

.6)

22

.0(2

1.8

)

[Ni(

L3S

L3)(

H2O

)]n

C10H

22N

6O

8N

i4

13

Red

[3

00

28

68

2.5

22

9.1

(28

.9)

5.3

(5.3

)2

0.3

(20

.2)


123

nitrogen atoms is expected to the reduce the electron

density in the azomethine link and lower the t(C=N)

absorption frequency. The peak due to t(C=N) are shifted

to lower frequencies and appears between 1633 and

1604 cm-1, indicating coordination of the azomethine

nitrogen to the nickel, cobalt or copper metal [29]. How-

ever, the absence of t(O–H) stretching bands in the i.r.

spectra of the complexes together with the existence of H-

bridge (O–H��O) at near 1731–1722 cm-1 and the shifting

of -C=N and -N–O stretches provide support for [MN4]n or

[MN4�H2O)]n-type coordinations in the metal complexes

[30]. A sharp band observed between 989 and 929 cm-1

for both the ligands and their complexes is assigned to the

t(N–O) vibration. The shifts of N–O absorbances to lower

energy are consistent with the protonation occurring at the

hydrogen-bridged oxime oxygen atoms, yielding a covalent

O–H bond. The formation of O–H bonds results in the

removal of electron density from N–O bond and

Table 2 Characteristic i.r.

bands (cm-1) of the ligands and

poly-metal complexes as KBr

pellets

Compounds O–H/N–H Aliph. C–H O–H��O N–O C=N M-N

Ligand (L1SL1H4) 3605–3060 2924–2870 – 986 1640 –

[Co(L1SL1)(H2O)]n 3623–3108 2922–2859 1724 944 1609 453

[Cu(L1SL1)]n 3322 2924–2871 1722 956 1613 495

[Ni(L1SL1)]n 3336 2918–283 1731 952 1605 460

Ligand (L2SL2H4) 3629–3084 2945–2872 – 982 1641 –

[Co(L2SL2)(H2O)]n 3647–3078 2959–2865 1722 942 1614 474

[Cu(L2SL2)]n 3351 2957–2871 1723 970 1615 480

[Ni(L2SL2)]n 3334 2918–2855 1727 953 1604 461

Ligand (L3SL3H4) 3546–3149 2965–2869 – 949 1647 –

[Co(L3SL3)(H2O)]n 3611–3125 2936–2865 – 932 1618 476

[Cu(L3SL3)]n 3323 2918–2866 1722 929 1630 470

[Ni(L3SL3)(H2O)]n 3365 2919–2871 1731 930 1633 479

Table 3 Characteristic u.v.-vis

bands of the ligands and poly-

metal complexes

*Shoulder

Compounds Solvents Wave length (kmax. (nm))

Ligand (L1SL1H4) C2H5OH 216 259 290

CHCl3 220 253 266*

[Co(L1SL1)(H2O)]n C2H5OH 257 312* 483 515 690

CHCl3 218 252 314* 428 634

[Cu(L1SL1)]n C2H5OH 221 258 294

CHCl3 215 256 310* 614

[Ni(L1SL1)]n C2H5OH 227 255 313 469 594

CHCl3 222 252 316* 460 549

Ligand (L2SL2H4) C2H5OH 216 260 297

CHCl3 215 257 272*

[Co(L2SL2)(H2O)]n C2H5OH 253 319* 467 558 685

CHCl3 215 244 422 470* 679

[Cu(L2SL2)]n C2H5OH 246 289 297 469 678

CHCl3 213 238 299* 456* 685

[Ni(L2SL2)]n C2H5OH 251 281* 416 489 540

CHCl3 218 280 304 477 558

Ligand (L3SL3H4) C2H5OH 216 260 294 310

CHCl3 216, 240, 257*

[Co(L3SL3)(H2O)]n C2H5OH 257, 301* 558, 661

CHCl3 213 236 276

[Cu(L3SL3)]n C2H5OH 250 257 303* 315

CHCl3 218 237 258

[Ni(L3SL3)(H2O)]n C2H5OH 258 306 361 486 575

CHCl3 239 316* 373 448* 560


123

corresponding increase in the N–O bond lengths and

decreased N–O stretching vibrations [31].

UV-v.i.s. Spectra

Electronic spectra of ligands (L1SL1H4, L2SL2H4, and

L3SL3H4) and their Ni(II), Cu(II), and Co(II) poly-metal

complexes have been recorded in the 200–900 nm range in

C2H5OH and CHCl3 solutions as shown in Table 1. Each

ligand and complex shows several intense absorptions in

the visible and ultraviolet region. The uv-vis spectra of the

ligands and metal complexes in C2H5OH showed three–

five absorption bands between 216 and 690 nm and in

CHCl3 showed three–five absorption bands between 213

and 685 nm, respectively. The wavelengths of absorption

bands in both solvents are practically identical. The bands

below 416 nm in C2H5OH or CHCl3 and are almost cer-

tainly associated with intraligand p ? p* and n? p* or

charge-transfer transitions [32]. In the electronic spectra of

the ligands and their Ni(II), Cu(II), and Co(II) metal

complexes, the wide a range bands seems to be due to both

the p ? p*, n? p* and d–d transitions of C=N and charge-

transfer transition arising from p electron interactions

between the metal and ligand, which involves either a

metal-to-ligand or ligand-to-metal electron transfer [33,

34]. The absorption bands below 297 nm in C2H5OH and

299 nm in CHCl3 are practically identical and can be

attributed to p ? p* transitions in the benzene ring or

azomethine (-C=N) groups. The absorption bands observed

within the range of 301–319 nm in C2H5OH and the range

of 304–316 nm in CHCl3 are most probably due to the

transition of n? p* of imine group corresponding to the

ligands or complexes [35]. Also, the absorption bands at

428–489 nm (in C2H5OH or in CHCl3 solvents) are

assigned to M?L charge transfer (MLCT) or L?M charge

transfer (LMCT) and 1A1g ?1B1g transitions [36],

respectively. The electronic spectrum of Cu(II) and Co(II)

in various solvents is significantly different from that in

C2H5OH or CHCl3, and shows a broad bands at range 515–

690 nm assigned to 2Eg ? 2T2g transitions, characteristic

for tetragonally, elongated octahedral or square planar

geometry [37, 38]. The weak d–d transitions of square

planar Ni(II) complexes were observed in the between 469

and 594 nm [39].

Magnetic moments

Magnetic moments measurements of the Ni(II) poly-metal

complexes carried out at 25 �C show that [Ni(L1SL1)]n and

[Ni(L2SL2)]n complexes are diamagnetic, indicating the

low-spin (S=0) square planar d8-systems, whereas

[Ni(L3SL3)(H2O)]n complex is paramagnetic (leff = 2.52

BM), indicating the the high-spin (S=1) octahedral d8-

systems, as expected. Since the [Ni(L1SL1)]n and

[Ni(L2SL2)]n metal complexes are diamagnetic, their n.m.r.

spectra could be obtained. The absorption bands observed

for the electronic spectra of Ni(II) complexes also support

the square planar or octahedral geometry for Ni(II) com-

plexes [40] (Table 3). The magnetic moments of the Cu(II)

complexes (Table 1) at room temperature are between 1.79

and 1.84 BM, which are typical for mononuclear Cu(II)

complexes with a S=1/2 spin-state and do not indicate

antiferromagnetic coupling of spins at this temperature.

The magnetic moments of the Co(II) complexes at room

temperature are also between 3.82 and 3.96 BM which are

close to the spin-only magnetic moments for three unpaired

electrons. The Co(II) complexes were characterized by i.r.

specta and elemental analyses and it was found that two

molecules of water were axially coordinated to the cobalt

ion. The observation of an (O–H��O) bond leads us to

Table 4 Voltammetric data for the metal complexes in DMSO-TBAP

Complexes M(III)/M(II) E1/2 (V) DEb(V) Ipc/Icpa M(II)/M(I) Epc(V) M(II)/M(I) Epa (V) DEb(V)

[Co(L1SL1)(H2O)]n -1.16 -0.16

[Co(L2SL2)(H2O)]n -1.09 -0.16

[Co(L3SL3)(H2O)]n -1.03 -0.20

[Cu(L1SL1)]n 0.36 0.23 0.50 -0.68 -0.31 0.37

[Cu(L2SL2)]n 0.38 0.24 0.46 -0.66 -0.30 0.36

[Cu(L3SL3)]n 0.39 0.25 0.48 -0.65 -0.29 0.36

[Ni(L1SL1)]n 0.66a -1.39

[Ni(L2SL2)]n 0.69a -1.30

a anodic peak potentials for oxidation of [Ni(L1SL1)]n and [Ni(L2SL2)]n

Epa: anodic peak potential for oxidation, Epc: cathodic peak potential for reduction for irreversible processesb DE = Epc - Epa at 0.100 V s-1 scan ratec ipc/ipa for the oxidation of the copper complexes at 0.100 V s-1 scan rate


123

consider the geometry of complexes to be octahedral [41]

(shown in Tables 1 and 2). The structures of the Ni(II),

Cu(II), and Co(II) poly-metal complexes proposed are

shown in Fig. 1.

Solubility and molar conductivity

The ligands (L1SL1H4, L2SL2H4 and L3SL3H4) are soluble

in THF, C2H5OH, CH3OH, CH2Cl2, DMSO, and DMF

solvents, whereas Ni(II), Cu(II), and Co(II) poly-metal

complexes are sparingly soluble in C2H5OH but are soluble

in DMSO and in DMF solvents to give stable solutions at

room temperature. With a view to studying the electrolytic

nature of the Ni(II), Co(II) and Cu(II) poly-metal com-

plexes, their molar conductivities were measured in DMF

at 10-3 M. The molar conductivities (KM) values of these

Ni(II), Co(II), and Cu(II) metal complexes are 23–53 X-

1cm2mol-1 at room temperature [42] indicating their

almost non-electrolytic nature. Due to non-free ions in

Ni(II), Cu(II), and Co(II) complexes, the results indicate

that Ni(II), Cu(II), and Co(II) metal complexes are very

poor in molar conductivity. The conductivities of the Cu(II)

complexes are rather higher than those of the other com-

plexes, perhaps due to electronic effects [43].

Electrochemistry

The electrochemical behaviours of the new vic-dioxime

complexes, [M(L1SL1)]n, [M(L2SL2)]n, and

[M(L3SL3)�H2O]n, where M = Co(II), Cu(II) and Ni(II)

ions, were investigated using cyclic voltammetric (CV) and

differential pulse voltammetric (DPV) techniques in

DMSO solution containing 0.1 M TBAP. The data

obtained in this work are summarized in Table 4. Figure 2a

shows the CV of the [Co(L3SL3)(H2O)]n complex where

the inset figure also exhibits DPV in the same experimental

condition. As seen, [Co(L3SL3)(H2O)]n displayed an irre-

versible one-electron reduction at Epc = - 1.03 V versus

Ag/AgCl at 0.100 V�s-1 scan rate. The complex also gave

an anodic peak at Epc = - 0.20 V upon reverse scanning.

This large peak separation (DE = 0.83 V) between the

reduction and its re-oxidation wave is probably due to the

coupled chemistry following the electrode reaction. The

reduction wave is assigned to metal-based character since

the reduction based on the Co(II)/Co(I) process, in vic-

dioxime complexes, are usually observed at more positive

values compared to that of ligand-based [24, 43, 44–47].

The anodic peak observed at Epc = - 0.20 V indicates

that the six-coordination of the cobalt complex

[Co(L3SL3)(H2O)]n changes into the four-coordination by

loosing its axial ligands following the reduction process

(Co(II)/Co(I)). On scanning through a second cycle no

change on the cathodic and anodic waves is observed for

these electrode processes; reductions occur at the same

potentials as observed in the first cycle. Multiple scan

resulted in nearly superimposable CV’s, thereby showing

the marked stability of the two reduction states of cobalt

involved in the electrochemical study. The inset figure

corresponding to the DPV of [Co(L3SL3)(H2O)]n in Fig. 2a

confirms the reduction process based on the Co(II)/Co(I)

couple. The electrochemistry of the [Co(L1SL1)(H2O)]n

and [Co(L2SL2)(H2O)]n complexes is similar to that of

[Co(L2SL2)(H2O)]n in the same experimental condition.

One difference is that the re-oxidation process of the

[Co(L2SL2)(H2O)]n is adsorption controlled on the elec-

trode surface, to which the proportional increasing of

anodic peak currents with increasing scan rates provides

satisfactory evidence for the adsorption character of wave

(Fig. 2b, Co(I)/Co(II) couple). The cathodic peak potential

-1,5 -1,0 -0,5 0,0 0,5

-1,5 -1,0 -0,5 0,0

-10

-5

0

5[Co(L3SL3)(H

2O)]

nCo(I)/Co(II)

Co(II)/Co(I)

/tnerru

CµA

Potential / V vs. Ag/AgCl

-1.5 -1.0 -0.5 0.0

-2

-1

0

Co(II)/Co(I)

/tnerr

uC

µA


-50

0

50

IB

[Co(L2SL2)(H2O)]

n

Co(II)/Co(I)

Co(I)/Co(II)

/tnerru

CµA


(a)

(b)

Fig. 2 (a) CV of [Co(L3SL3)(H2O)]n at 0.100 V�s-1 scan rate, where

the inset figure shows DPV of the complex; (b) CVs of

[Co(L2SL2)(H2O)]n at the scan rates of 0.025 - 0.250 V�s-1 in

DMSO/0.1M TBAP. IB represents the CV of 0.1 M TBAP in DMSO

without the complexes


123

of [Co(L1SL1)(H2O)]n is negatively shifted by 0.07 V

compared to that of the [Co(L2SL2)(H2O)]n due to the

presence of the phenyl and p-tolyl substituents on the

oxime-moieties, both of which are electron-withdrawing

groups and their effects decrease in the given order. On the

other hand, the cathodic peak potentials of [Co(L1SL1)(-

H2O)]n and [Co(L2SL2)(H2O)]n appeared at the lower

reduction potentials in respect to that of [Co(L3SL3)(-

H2O)]n on contrast to the expectation, probably due to the

bulky group effect of the polymeric structure.

The copper complexes [Cu(L1SL1)]n, [Cu(L2SL2)]n, and

[Cu(L3SL3)]n showed one oxidation and one reduction

processes and their cathodic peak potentials of the reduc-

tion processes are at Epc = - 0.68, -0.66, and -0.65 V

and their half-wave peak potentials of the oxidation pro-

cesses at E1/2 = 0.36, 0.38, and 0.39 V, respectively versus

Ag/AgCl at 0.100 V�s-1 scan rate (Table 4). An example

of the CVs and DPV of the copper complexes (for the

[Cu(L2SL2)]n complex) is given in Fig. 3. As seen, the

complex exhibits a quasi-reversible oxidation process in

the scan rates of 0.025–0.250 V�s-1 with cathodic–anodic

peak separations (DE = 0.191 - 0.347 V) and the ratio

of anodic to cathodic peak currents (ipa/ipc = 0.46). The

deviation from the reversibility to the quasi-reversibility is

probably due to the decomposition of the electro-oxidized

species to small extent on the electrode surface in the time

scale of the CV measurement. The oxidation process is the

metal-based Cu(II)/Cu(I) couples as previously observed

for some vic-dioxime copper complexes [24, 43, 44–47].

The copper complex also exhibits an irreversible reduction

process in the scan rates of 0.025–0.250 V�s-1. The com-

parative electrochemical study between cobalt and copper

complexes indicates that the shifts of the cathodic or half-

wave peak potentials observed between the copper

complexes are similar to what are observed between the

cobalt complexes. An example of the CV of the nickel

complexes is depicted in Fig. 4. As seen, the nickel com-

plexes show irreversible one reduction and one oxidation

processes based on the metal. Their cathodic and anodic

peak potentials are shifted to the higher potentials com-

pared to those of the cobalt and copper complexes,

indicating the harder reducibility or oxidizibility of the

nickel complexes with respect to the copper and cobalt

complexes in the same experimental condition.

Acknowledgments This work have been supported, in part, by the

Research Fund of Harran University (Sanliurfa, Turkey). This work

has also been supported, in part, by the Turkish Academy of Sciences

in the framework of the Young Scientist Award Program (TUBA-

GEB_IP).

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novel vic-dioxime ligands and their poly-metal complexes bearing 1,8-diamino-3,6-dioxaoctane:...

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