schiff base complexes

Binuclear copper(II), nickel(II) and cobalt(II) complexes with N2O2

chromophores of glycylglycine Schiff-bases of acetylacetone,

benzoylacetone and thenoyltrifluoroacetone

Shehab A. SallamChemistry Department, Faculty of Science, Suez Canal University, Ismailia, Egypt

Received 9 June 2005; accepted 10 August 2005

Abstract

Binuclear copper, nickel and cobalt complexes of the Schiff-bases obtained by condensation of glycylglycine withacetylacetone, benzoylacetone, dibenzoylmethane and thenoyltrifluoroacetone were prepared by template synthesis.The complexes were characterized by elemental analysis, conductivity measurements, magnetic moments, i.r.,u.v.�vis. spectra, e.s.r., X-ray diffraction, t.g.a., d.t.a. and d.s.c. thermal analysis. All the complexes are non-electrolytes with low magnetic moments that indicate spin�spin or antiferromagnetic exchange interactions.Spectral properties support square planar and square pyramidal or trigonal bipyramidal structure provided by theN2O2 chromophores. E.s.r. spectra of the copper complex confirm the binuclear structure and the presence ofmagnetic interaction. Thermal studies supported the chemical formulation of these complexes and showed that theydecompose in three to four steps depending on the type of ligand. Activation energies Ea and enthalpies DH,associated with the thermal decomposition of the complexes were calculated and correlated with the type ofcomplexed metal. A mechanism for thermal decomposition is proposed for the complexes.

Introduction

Systems in which the Schiff-bases of amino acids arederived from b-diketones have received scant interest[1�3]. Neglect of this ligand type of b-ketoimine-N-al-kanoic acid seems to arise from their known hydro-lytic instability [4]. Furthermore, those metalcomplexes primarily have been confined to transitionelements. N-benzoylacetoneglycine has been stabilizedin the form of its sodium salt and the complexes ofthis ligand with a series of lanthanide(III) ions havebeen prepared and their antioxidative actions alsohave been determined [5].However, little attention has been paid to systems in

which the Schiff-bases are derived from simple peptides.A vanadium complex VO(sal�glygly)(H2O)n (sal�gly-gly = N-salicylideneglycylglycine); n = 1.5�3.0 hasbeen isolated from relatively concentrated solutionscontaining oxovandium(IV), glycylglycine and salicylal-dehyde [6]. Synthesis, crystal structure and magneticstudies of cis-configuration copper(II)-M(II) (M = Ba,Ca) complexes of the sal�glygly Schiff-base weredetermined [7]. We have prepared and characterized theuranyl complexes of Schiff-bases obtained by condens-ing glygly with hydroxybenzaldehydes andhydroxynaphthaldehyde [8]. Recently, two copper(II)tripeptide Schiff-base complexes: [Mg(H2O)6] [CuL]2 Æ3.5H2O and [Cd(H2O)4(CuL)2] Æ 3 . 5H2O (H3L =

N-sal�glygly) have been synthesized and structurallycharacterized [9]. Also, reaction of Na[CuL] (whereH3L = the Schiff-base derived from 5-bromosalicylal-dehyde and glylgly) with La(NO3)3 leads to theformation of carboxylate-bridged polynuclear cop-per(II)�lanthanide(III) rings of formula [Ln(H2O)5(CuL)2][CuL] Æ 8H2O (Ln = La and Ce) [10].In continuation of our interest in the coordination

behavior of Schiff-bases of the biologically importantglycylglycine, we report the synthesis and characteriza-tion of a series of copper(II), nickel(II) and cobalt(II)complexes of Schiff-bases obtained by condensingglycylglycine (glygly) with acetylacetone (acac), ben-zoylacetone (ba), dibenzoylmethane (dbm) and the-noyltrifluoroacetone (tta). The following Schiff-baseswere synthesized in solution: L1(glygly + acac), L2(gly-gly + ba), L3(glygly + dbm) and L4(glygly + tta).The complexes were characterized using: elementalanalysis; conductivity measurements, magnetic momentdeterminations; X-ray diffraction; e.s.r., u.v.�vis., i.r.,and thermal analysis (t.g., d.t.g., d.t.a. and d.s.c.).

Experimental

Materials

Glycylglycine was obtained from BDH. Acetylacetone, benzoylacetone, dibenzoylmethane and then-E-mail: [email protected]

Transition Metal Chemistry (2006) 31:46�55 � Springer 2006DOI 10.1007/s11243-005-6312-4

oyltrifluoroacetone were purchased from Fluka andwere used as supplied. Other chemicals are reagentgrade and were used without further purification.

Template synthesis of the complexes

To a solution of glycylglycine (0.324 g, 2 mmol) andsodium acetate (0.328 g, 4 mmol) in H2O (10 cm3), asolution of the b-diketone (1 mmol) in absolute EtOH(5 cm3) was added dropwise with stirring. The reactionmixture was heated with stirring on a water bath for2 h during which time the color changed to yellow,indicating formation of the corresponding Schiff-bases.The respective metal chloride (2 mmol) in absoluteEtOH (5 cm3) was added dropwise to the reactionmixture. The whole mixture was boiled with stirringfor 30 min. The precipitated complexes were then fil-tered, washed several times with EtOH and Et2O, anddried in vacuum over anhydrous CaCl2.

Physical measurements

C, H, and N were estimated using a Heraus CHN-rapid analyzer. Metal analyses were carried out on anARL 3410 ICP sequential spectrophotometer at wave-lengths 324.75, 231.60 and 228.62 nm for Cu, Ni andCo, respectively. The i.r. spectra were recorded (KBrdisc) in the 4000�400 cm)1 range on Nicolet Impact400 and Bruker Vector 22 spectrophotometers. Theelectronic absorption spectra were obtained using nu-jol of mull and 10)3 M DMSO solutions in a 1 cm cellusing a PerkinElmer S550 spectrophotometer. Solidstate x-band e.s.r. spectra were recorded on a VarianE4 spectrophotometer. Magnetic susceptibility mea-surements were carried out using the modified Gouymethod on a Johnson�Matthey balance at room tem-perature and using mercury(II) tetrathiocyanatocobalt-ate(II) as the calibrant. The effective magneticmoments, leff, per metal atom was calculated from theexpression:

�eff ¼ 2:83ffiffiffiffiffiffiffiffi

�:Tp

B.M.

where v is the molar susceptibility corrected using Pas-cal’s constant for the diamagnetism of all atoms in the

complexes. X-ray powder diffraction was performedusing a Shimadzu XD-3 diffractometer with Cu-Ka

radiation. T.g.a., d.t.g., d.t.a. and d.s.c. were recordedon a Perkin�Elmer 7 Series thermal analyzer equippedwith Pyres software under a dynamic flow of nitrogen(10 L/min.) and heating rate of 10�C/min from ambi-ent temperature to 1000�C (t.g.a and d.t.a.) or to600�C (d.s.c.). The number of decomposition steps wasidentified using d.t.g. Electrical conductivity measure-ments were carried out at room temperature on freshlyprepared 10)3 M DMSO solutions using a WTW con-ductivity meter fitted with a L100 conductivity cell.

Results and discussion

All the metal complexes are insoluble in common or-ganic solvents such as methanol, ethanol, benzene,chloroform, acetone, dichloroethane and diethyl ether,but partly soluble in DMSO. The analytical data(Table 1) show that the complexes can be representedas [M2 Æ L1] Æ 2H2O where M = CuII, NiII and CoII

and [M2 Æ L2�4 �Cl2 Æ (H2O)2] Æ nH2O where M = NiII

and CoII.The low molar conductance values of 10)3 M solu-

tions of the complexes in DMSO at room temperaturelie in the 17�55 lS range which show that they are allnon-electrolytes [11] and that the anion is also coord-inated to the metal ion.

Infra-red spectra

Free glycylglycine has a strong absorption band at3300 cm)1, which corresponds to the NH2 stretchingfrequency [12]. In addition, a weak band near3080 cm)1 was noted and assigned tentatively to anadditional NH absorption [12]. Uncoordinated acetyl-acetone contains the free carbonyl entity that has anintense absorption band at ca. 1700 cm)1 [13]. Allthese bands are not observed in the spectra of thecomplexes [Cu2 Æ L1] Æ 2H2O, [Ni2 Æ L1] Æ 2H2O, and[Co2 Æ L1] Æ 2H2O (Table 2). Besides, the strong bandsaround 1530 and 1280 cm)1, which are characteristicof stretching and bending modes of the bonded

Table 1. Analytical data, electrical conductivities and magnetic moments of the synthesized complexes

Complex Color Mol. wt. M.p.

(�C)Found (calcd.)% X (lS) leff B.M.

C N H M

[Cu2 Æ L1] Æ 2H2O Green 487.13 246 32.4(32.0) 11.1(11.5) 3.8(4.1) 26.0(26.1) 17 0.31

½Ni2 � L1� � 2H2O Olive 477.43 345 32.5(32.7) 11.1(11.7) 4.0(4.2) 25.3 (24.6) 28 Dia.

[Co2 Æ L1] Æ 2H2O Pink 477.91 295 32.8(32.7) 11.3(11.7) 3.8(4.2) 25.1(24.7) 47 1.25

[Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O Olive 515.35 207 32.9(32.6) 5.5(5.4) 3.43(3.7) 22.65(22.8) 33 0.6

[Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2 H2O Pink 533.82 300 chr 31.6 (31.5) 5.4 (5.2) 3.6(3.55) 22.2 (22.1) 45 0.74

[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O Olive 613.38 360 37.3(37.2) 4.7(4.6) 4.0(4.1) 19.3(19.3) 31 0.80

[Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O Pink 613.86 >400 37.2 (37.2) 4.7 (4.6) 3.85(4.1) 19.34(19.2) 43 0.64

[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2 1/2 H2O Pink 602.82 307 chr 23.8(23.9) 4.3(4.6) 3.01(2.8) 19.8 (19.55) 55 1.13

chr=charring.

47

carbonyl of acetylacetone, are absent [13]. The abovei.r. data clearly indicate that, in these complexes, thecarbonyl groups of one molecule of acetylacetone havereacted with the amine groups of two molecules of gly-cylglycine through the template effect of the metalatoms. This contention is supported by the presence ofnew bands in the spectra of the complexes in the2927�2965, 1321�1330, 1267�1275, 1103�1130 and609�615 cm)1 ranges which are characteristic of theacetylacetone moiety and may be assigned to �CH3,dsCH3, �C�CH3, dCH + �C�CH3 and to ring defor-mation, respectively [13�15]. Also, new bandscharacteristic of the �(NCO) appear in the 1515�1549and 1321�1327 ranges due to coordination throughimidol oxygen [16]. I.r. spectra of the complexes[Ni2 Æ L2 Æ Cl2] Æ H2O, [Co2 Æ L2 Æ Cl2] Æ 2H2O, [Ni2 Æ L3

Æ Cl2] Æ 3H2O, [Co2 Æ L3 Æ Cl2] Æ 3H2O and [Co2 Æ L4

Æ Cl2 Æ (H2O)2] Æ 2 1/2H2O show the disappearance of�(C ¼ O) of the b-diketone part and the appearance ofnew bands in the 1518�1567 cm)1 range assignable to�(C�O) which suggest that the carbonyl group is in-volved in coordination in the enol form through de-protonation [17].The i.r. spectra of all the complexes show broad

bands in the 3332�3368 cm)1 range that may be wellcorrespond to �N�H amide. Also, a sharp strong bandin the 1637�1646 cm)1 range indicates the presence ofbonded azomethine groups (>C ¼ N) [17]. The strongbroad bands, present in the 1575�1587 cm)1 range inthe complexes’ spectra, are assigned to an asymmetricstretch of coordinated carboxylate groups. Bandscorresponding to �sym(COO)) appear in the1429�1444 cm)1 range, giving a frequency difference(D� value) of 131�158 cm)1 compared with 171 cm)1

of free glygly that reflect the monodentate nature ofthe carboxylate groups [18, 19]. Strong absorption

bands in the 1398�1405 cm)1 range could be assignedto the C�N amide stretching mode [16]. The carboxylwagging vibrations appear in the 755�760 cm)1 rangein the complexes’ spectra. Bands in the 411�420 cm)1

range in the complexes’ spectra are assigned to theM�O stretch. The �(M�N) vibration is attributed tothe bands in the 443�525 cm)1 range [13].

Magnetic and spectral properties

The magnetic moments of the complexes calculatedfrom the corrected magnetic susceptibility determinedat room temperature are given in Table 1. The mag-netic moments reported for the complexes are lowerthan the spin only values. The subnormal magneticmoments observed for the Cu, Ni, and Co complexesare accounted for by assuming a binuclear structurewith considerable interaction between metal ion sys-tems owing to spin�spin coupling causing an antifer-romagnetic interaction mechanism to operate throughthe orbitals of the metal ions [20�22].The magnetic moment per copper atom in the

[Cu2 Æ L1] Æ 2H2O complex is below the ‘spin only’value of 1.73 B.M. for copper(II). Such lowering ofthe magnetic moment has been observed in the bi- andtrinuclear complexes, which contained only copper(II)and has been shown to arise from antiferromagneticexchange interactions [20, 23�25]. The observed mag-netic moment indicates fairly strong magneticexchange interaction between pairs of neighboringcopper atoms, a phenomenon which occurs in a largenumber of copper(II) complexes [20].The copper complex possesses square planar geome-

try as evidenced by the appearance of only one bandin the electronic spectra (nujol mull) at 14306 cm)1

with two shoulders at 18135 and 22988 cm)1. These

Table 2. Spectral data of the complexes

Complex I.r. spectral data. (cm)1) Vis. spectral data (cm)1)

mOH mNH mC=N masCOO) mC�O msymCOO) mC�N dCOO) M�OM�N

Nujol DMF

[Cu2 Æ L1] Æ 2 H2O 3400 s. sh 3232 s. sh 1637 s 1587 s 1515 m 1429 ssh 1403 m 760 m 411 m 14,204 sh 14,306 br

3082 s. sh 484 m 18,135 11,152 sh

22,988

½Ni2 � L1� � 2H2O 3430 m. sh 3361 s. m 1637 s 1587 s � 1429 ssh 1405m 760 m 420 m 14,925 13,458

3273 s. sh 484 m 17,699 14,749

25,641 23,923

[Co2 Æ L1] Æ 2H2O 3451 s. sh 3366 s. sh 1641 s 1585 s � 1444 m sh 1403 m 755 m 443 m 18,281 br 15,128

3287 s. sh 525 m 16,051 sh

[Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O 3451 m. br 3064 m 1600 s 1564 s 1518 s 1454 m sh 1398 br 763 m 428 m 13,020 sh 13,254 sh

452 m 15,748 br 15,220 br

[Co2 Æ L2 Æ Cl2 Æ H2O] Æ 2H2O 3451 m. br 3366 s. sh 1646 s 1575 s 1552 s 1444 m sh 1400 m 759 m 435 m 18,050 br 19,607

3268 m. br 464 m

[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 3454 m. sh 3366 s. sh 1646 s 1575 s 1567 s 1444 m sh 1406 m 759 m 417 m 13,254 sh 13,054

3281 m. sh 467 m 15,037 br 14,641

[Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 3458 m. br 33656 m. b 1640 s 1582 s 1555 s 1441 m sh 1405 m 760 m 417 m 17,960 br 19,540

3268 m. br 462 m

[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2 1/2H2O 3462 m. br 3363 m. br 1646 s 1578 s 1552 s 1444 m sh 1400 m 759 m 412 m 17,985 br 19,230

3279 m. br 467 m

48

bands could be assigned to 2B1g fi 2A1g,2B1g fi 2B2g

and 2B1g fi 2Eg transitions, respectively [26]. Thesquare planar coordination of each copper atom isprovided by the two N2O2 chromophores of the Schiff-base. One of the copper atoms forms a six-memberedring and two five-membered rings. The second copperatom forms two five-membered rings. This geometrymay probably lead to a greater antiferromagneticinteraction [27]. Upon dissolution in DMSO, the com-plex changes from green to violet and the spectra ofthe solution show a broad band at 14306 cm)1 with ashoulder at 11152 cm)1. This indicates that, in solu-tion, the copper complex changes geometry fromsquare planar to distorted octahedral with 2Eg fi 2T2g

and 2B1g fi 2A1g transitions through the adduction ofsolvent molecules [28].The complex [Ni2 Æ L1] Æ 2H2O displays diamagnetic

behavior. This is consistent with square-planar geome-try of the two N2O2 chromophores around the Ni(II).This interpretation is supported by electronic spectrain the nujol mull (Table 2). It shows three absorptionbands at 14925, 17699 and 25641 cm)1 which areattributed to 1A1g fi 1A2g(�1),

1A1g fi 1B1g(�2) and1A1g fi 1Eg(�3) transitions [29]. A spectrum of thecomplex in DMSO shows bands at 13458, 14749 and23923 cm)1 respectively, associated with a colorchange from olive to red in solution. This indicates achange of nickel coordination from four (square pla-nar) to six (Oh) with 3A2g(F) fi 3T2g(F)(�1),3A2g(F) fi 3T1g(F)(�2) and 3A2g(F) fi 3T1g(P)(�3)transitions [29].The leff values per nickel(II) ion for the complexes

[Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O and [Ni2 Æ L3 Æ Cl2 Æ(H2O)2] Æ 3H2O are 0.6 and 0.8 B.M. which differ fromthe value found for the complex [Ni2 Æ L1] Æ 2H2O. Thiscould indicate that the coordination geometry is differ-ent in the two complexes compared with the latter one.Nujol mull electronic spectra of the two complexesshowed a broad band in the 15037�15748 cm)1 rangewith a shoulder in the 13020�13254 cm)1 range. Thissuggests distorted trigonal bipyramidal or square pyra-midal geometry around the nickel(II) ion [30]. Thesame absorption peaks are found when the spectra, ob-tained for DMSO solution of the complexes, dissolvewithout any color change. This indicates preservationof the five coordinate configuration in donor solvents.Magnetic moment values for the cobalt(II) complex

showed that there is also an antiferromagnetic exchangeinteraction present. The complex [Co2 Æ L1] Æ 2H2Oexhibits a broad band at 18281 cm)1 which is ascribedto the 1A1g fi 1B1g transition as observed in the case ofsquare-planar cobalt complexes [28]. Solution of thecomplex in DMSO induces the following color changes:pink fi violet fi blue fi green. The electronic solu-tion spectra were characterized by the appearance oftwo main bands at 15128 and 16051 cm)1 which can beascribed to the 4T1g fi 4A2g(F)(�2) and 4T1g fi4T1g(P)(�3) transitions of an octahedral geometry.

The d�d band appearing in the absorption spectra(nujol mull) of the complexes [Co2 Æ L2 Æ Cl2 Æ(H2O)2] Æ 2H2O, [Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O and[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2 1/2H2O were found tohave broad bands at 18050, 17960 and 17985 cm)1

respectively. Such transitions are characteristic ofsquare-pyramidal or distorted trigonal-bipyramidalstructures with the 4A2(F) fi 4A2(P) transition [28].The complexes dissolve in DMSO with no colorchange and their absorption bands are shifted to19607, 19540 and 19230 cm)1 respectively.

E.s.r. spectrum

The x-band e.s.r. spectrum of the copper complex wasrecorded on a powder solid at room temperature (Fig-ure 1). It exhibits an axial signal which can be inter-preted in terms of tetrahedral species with a strongsignal in the low field region, corresponding tog? ¼ 2:19, and a weak signal in the high field regiondue to gk ¼ 2:06 while the giso is 2.15. Splitting of thesignal in the high field region may be due to a differ-ence in surroundings between the two copper(II) cen-ters suggesting a binuclear structure for this complex[31]. As shown before, one of them forms a six-mem-bered ring and two five-membered rings, whereas theother forms only two five-membered rings. The G val-ue is equal to 0.32, indicating that antiferromagneticinteraction is taking place between the copper(II) ions[32,33]. It is interesting to note that theg? > gk > 2:0023, indicating that the unpaired electronis present in the dxy ground state in a square planargeometry [34].

X-ray powder diffraction

X-ray patterns of the complexes [Cu2 Æ L1] Æ 2H2O,[Ni2 Æ L1] Æ 2H2O, and [Co2 Æ L1] Æ 2H2O were recordedover the 2h = 0�100 range. The principal intraspacelines (d) and their intensities are represented in Table 3and are shown in Figure 2. Generally, the complexes[Cu2 Æ L1] Æ 2H2O and [Ni2 Æ L1] Æ 2H2O have nearlysimilar crystallinity which is higher than that of

Fig. 1. X-band powder e.s.r. of the complex [Cu2 Æ L1] Æ 2H2O.

49

[Co2 Æ L1] Æ 2H2O. Each complex has specific d values,which can be used for its characterization [35].

Thermal analysis

The thermal behavior of the complexes has been stud-ied using t.g.a., d.t.g., d.t.a. and d.s.c. analysis. Typicalcurves are present in Figures 3 and 4. Thermal curvesobtained for most of the complexes were very similarin character. The decomposition stages, temperature

ranges, decomposition product, as well as the foundand calculated weight loss percentages of the com-plexes are given in Table 4.The complexes [Cu2 Æ L1] Æ 2H2O, [Ni2 Æ L1] Æ 2H2O

and [Co2 Æ L1] Æ 2H2O undergo decomposition mainlyin four stages. The first stage takes place in the60�117, 50�223 and 62�110�C ranges with d.t.g.peaks at 103, 161 and 92�C, respectively, correspond-ing to dehydration and release of two water molecules.The mass loss observed in this step is 6.62, 7.46 and7.53% against the calculated loss of 7.38, 7.53 and7.52%, respectively. The dehydration step is associatedwith endothermic changes in the 107�150 and98�166�C ranges with d.t.a. peaks at 135�C, and inthe 155�187 and 150�196�C ranges with d.s.c. peaksat 172�C for the complexes [Cu2 Æ L1] Æ 2H2O and[Co2 Æ L1] Æ 2H2O. Dehydration of the complex[Ni2 Æ L1] Æ 2H2O takes place in two steps as shown bythe d.t.a. curve (Figure 2) which depicts two endother-mic peaks at 96 and 185�C in the temperature range75�218�C. This indicates that the complex dehydratesin two well separated steps due to the presence of aloosely bound lattice water molecule and more strong-ly bound one. Also, the formation of hydrogen bondswith the lattice water may raise the dehydration tem-perature. The wide range of dehydration shown by thet.g. results (50�223�C) supports this idea. The d.s.c.curve of this complex shows an endothermic peak at

Table 3. X-ray powder diffraction of the [Cu2 Æ L1] Æ 2H2O, [Ni2 Æ L1] Æ 2H2O and [Co2 Æ L1] Æ 2H2O complexes

[Cu2 Æ L1] Æ 2H2O ½Ni2 � L1� � 2H2O [Co2 Æ L1] Æ 2H2O

n I 2h d n I 2h d n I 2h d

1 100 7.465 11.833 1 100 6.340 13.929 1 100 6.368 13.868

2 70 6.660 13.261 2 15 5.850 15.095 2 10 5.870 15.044

3 30 5.299 16.663 3 95 5.351 16.502 3 75 5.344 16.523

4 40 4.811 18.352 4 25 4.808 18.364 4 100 4.688 18.834

5 25 4.135 21.351 5 92 4.687 18.838 5 10 4.003 22.055

6 37 4.058 21.756 6 18 3.986 22.149 6 20 3.712 23.783

7 60 3.750 23.542 7 58 3.690 23.925 7 77 3.563 24.777

8 85 3.661 24.114 8 30 3.631 24.314 8 10 2.859 30.877

9 20 3.336 26.463 9 70 3.553 24.847 9 12 1.820 48.501

10 10 3.213 27.476 10 50 3.498 25.238

11 12 3.049 28.953 11 32 2.589 34.096

12 27 2.790 31.640 12 25 2.348 37.595

13 10 1.810 48.769

Fig. 2. X-ray powder diffraction of the Cu(II), Ni(II) and Co(II)

complexes of the L1 Schiff-bases.

Fig. 3. T.g.a. and d.t.g. of the Cu(II), Ni(II) and Co(II) complexes of the L1 Schiff-bases.

50

239�C representing dehydration in the temperaturerange 217�262�C. The second decomposition stage ofthe complexes [Ni2 Æ L1] Æ H2O and [Co2 Æ L1] Æ H2Otakes place in the 280�328 and 224�308�C rangeswith d.t.g. peaks at 303 and 270�C. This step bringsweight losses of 16.2 and 16.05% against the calcu-lated losses of 16.12% and 16.11% that correlate withelimination of 1 mol of both CH3COOH and NH3.The d.t.a. curve of the complex [Ni2 Æ L1] Æ H2O has awide range of thermal change between 308 and 443 �Cwith two endothermic peaks at 334 and 400 �C, whichcorrelates with melting and partial decomposition.Melting of the complex is also confirmed by a sharpendothermic d.s.c. peak at 337 �C. Endothermic chan-ges were observed for the complex [Co2 Æ L1] Æ H2O inthe 276�322�C range with a d.t.a. peak at 293 �C, andin the 283�296 range with a d.s.c. peak at 283 �Cwhich accompanied partial decomposition of the com-plex. On the other hand, the complex [Cu2 Æ L1] Æ H2Oloses CH3COOH and NH3 in two consecutive steps inthe 152�200 and 200�220�C range with d.t.g. peaks at184 and 208�C and mass losses of 11.34 and 3.95%against the calculated losses of 12.31 and 3.48%. Thisstep is followed by a steady mass loss and may be dueto the expulsion of the remaining part of the ligandmolecule together with volatilization of the residue[36]. This is supported by the endothermic peaks at213�C (d.t.a.) and 217�C (d.s.c.) corresponding tomelting of the complex, followed by sharp exothermicpeaks at 240�C (d.t.a.) and 240�C (d.s.c.) which repre-sent vaporization following melting [37]. The thirddecomposition stage of the complexes [Ni2 Æ L1] Æ H2O

and [Co2 Æ L1] Æ H2O is in continuation with the secondstage in the 328�365 and 308�324�C range with d.t.g.peaks at 351 and 323 �C. The mass loss observed is12.27 and 12.29% against a calculated loss of 12.56and 12.55% showing that another molecule ofCH3COOH is expelled. The third decomposition stageof the complex [Ni2 Æ L1] Æ H2O is shown also by theendothermic peak at 400�C (d.t.a. and d.s.c.). The finaldecomposition of the complexes with metal formationoccurs in the 480�670, 507�677 and 450�652�C ran-ges with d.t.g. peaks at 563, 577 and 550 �C, respec-tively.The t.g. curves of the complexes [Ni2 Æ L2 Æ Cl2 Æ

(H2O)2] Æ H2O and [Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2Oshow a weight loss of 3.5 and 8.96% (calcd. 3.49 and8.79%) in the temperature ranges 107�145 and122�223�C with d.t.g. maxima at 126 and 184 �C, cor-responding to a loss of lattice water. The seconddecomposition step of the two complexes in the220�290 and 310�394�C range have d.t.g. peaks at256 and 352 �C. It brings a weight loss of 32.11 and28.54% (calcd. 32.38 and 27.18%) which correlate withthe loss of two molecules of coordinated water, onechlorine molecule and partial decomposition of the li-gand with evolution of the carboxylate moiety as aceticacid. Partial decomposition of the complexes is contin-uing in the 290�390 and 394�496�C ranges with d.t.g.peaks at 348 and 445 �C showing weight losses of 21.3and 11.25% (calcd. 21.34 and 11.41%). The remainingpart of the ligands is lost in the 390�446 and496�594�C range where the final decomposition takesplace at 418 and 545 �C with the formation of nickelmetal as the final product.The endothermic peaks obtained in the 135�215 and

150�260�C range with d.t.a. peaks at 162 and 215�C forthe complexes [Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O and[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O are associated with thedehydration process. The d.s.c. curves of the complexeshave endothermic peaks at 156 and 258�C confirmingthe dehydration processes. Partial decomposition of thecomplex [Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O is indicated bythe endothermic peak in the 215�383�C range whichhas a maximum at 311 �C (d.t.a.). The complex[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O has endothermic chan-ges in the 318�368�C range with two d.t.a. peaks at 341and 400 �C. The first peak indicates melting of the com-plex while the second shows partial decomposition. Asharp endothermic peak in the 348�351�C range with amaximum at 349 �C (d.s.c.) supports melting of the com-plex, while partial decomposition is shown by the endo-thermic peak in the 386�405�C range with the d.s.c.maximum at 396�C. Final decomposition of the twocomplexes is shown by exothermic peaks at 551�C(d.t.a), 528�C (d.s.c.) for the complex [Ni2 Æ L2 Æ Cl2 Æ(H2O)2] Æ H2O and 554 �C (d.t.a.), 568 �C (d.s.c.) forthe complex [Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O.Thermal decomposition curves of the complexes

[Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O, [Co2 Æ L3 Æ Cl2 Æ(H2O)2] Æ 3H2O and [Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2½H2O

Fig. 4. D.t.a. calorigrams of the Cu(II), Ni(II) and Co(II) complexes

of the L1 Schiff-bases. (...) [Cu2 Æ L1] Æ 2H2O,(� Æ �)[Ni2 Æ L1] Æ2H2O,(�) [Co2 Æ L1] Æ 2H2O

51

each show a similar sequence of four decompositionsteps with elimination of the following species:1. Lattice water in the 82�114, 76�115 and 76�146�C

range with d.t.g. peaks at 98, 96 and 113�C withmass loss of 6.7, 8.32 and 7.5% respectively, (calcd.6.74, 8.79 and 7.46%).

2. Two molecules of coordinated water and one mole-cule of chlorine are evolved in the second decomposi-tion step of the complexes [Co2 Æ L2 Æ Cl2 Æ(H2O)2] Æ 2H2O and [Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2½-H2O. This step takes place in the 277�328 and271�400�C range with d.t.g. peaks at 302 and 335�C.It has a mass loss of 19.51 and 17.13% against a calcu-lated mass of 20.02 and 17.73%. The second decom-position step of the complex [Co2 Æ L3 Æ Cl2 Æ(H2O)2] Æ 3H2O represents the loss of two coordinatedwater molecules, one molecule of chlorine in additionto one molecule of acetic acid with a mass loss of27.88% (calcd. 27.18%). It starts at 247�C and endsat 301�Cwith a d.t.g. peak at 274�C.

3. Partial decomposition starts with the third step,which follows immediately after the second step inthe 328�342 and 400�520�C range, with d.t.g.peaks at 335 and 460�C. This step includes separa-tion of one molecule of acetic acid from thecomplexes [Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O and

[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2½H2O with a mass lossof 10.96 and 10.95% (calcd. 11.23 and 9.95%). Forthe complex [Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O, itsthird decomposition step is represented by partialdecomposition of the ligand in the 311�443�Crange with a d.t.g. peak at 377�C and mass loss of26.92% (calcd. 26.87).

4. Final decomposition takes place in the 342�521,450�615 and 520�787�C ranges with d.t.g. peaks at431, 541 and 653�C respectively.

The d.t.a. and d.s.c. curves of the complexes[Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O, [Co2 Æ L3 Æ Cl2 Æ(H2O)2] Æ 3H2O and [Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2½H2Oshow endothermic peaks within the temperature ranges93�195, 93�134 and 105�178�C, respectively. Thesepeaks have maxima at 148, 134 and 149�C which areassigned to dehydration. The d.s.c. curves also indicatedehydration with endothermic peaks at 193, 174 and198�C, respectively. The loss of coordinated water andchloride ions as chlorine gas are shown by the endo-thermic peaks at 301, 299�C (d.t.a.) and 316, 313�C(d.s.c.), for the complexes [Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ2H2O and [Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2½H2O. Thisfragment was lost in addition to decarboxylation asacetic acid � in one step � for the complex [Co2 ÆL3 Æ Cl2 Æ (H2O)2] Æ 3H2O as indicated by the endother-

Table 4. T.g.a. and d.t.g. data of the complexes

Complex T.g.

plateau

D.t.g.

(�C)Mass loss % Process Product Residues% and type

Exp. Calcd. Exp. Calcd.

[Cu2 Æ L1] Æ 2H2O 60�117 103 6.62 7.38 Dehydration 2H2O

152�200 184 11.34 12.31 Partial decomposition CH3COOH Not complete*

200�220 208 3.95 3.48 Partial decomposition NH3

480�670 563 � � Final decomposition L

½Ni2 � L1� � 2H2O 50�223 161 7.46 7.53 Dehydration 2H2O

280�328 303 16.2 16.12 Partial decomposition CH3COOH+NH3(0.23L)

328�365 351 12.27 12.56 Partial decomposition CH3COOH(0.18L) 23.01 24.58

507�677 577 41.06 39.18 Final decomposition 0.58L Ni

[Co2 Æ L1] Æ 2H2O 62�110 92 7.53 7.52 Dehydration 2H2O

224�308 270 16.05 16.11 Partial decomposition CH3COOH+HH3(0.23L)

308�324 323 12.29 12.55 Partial decomposition CH3COOH(0.18L) 25.16 24.66

450�652 550 39.29 39.14 Final decomposition 0.58L Co

[Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O 107�145 126 3.5 3.49 Dehydration 2H2O

220�290 256 32.11 32.38 Partial decomposition 2H2O+Cl2+CH3CO2H(0.22L)

290�390 348 21.3 21.34 Partial decomposition 0.4L 22.01 22.77


[Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O 82�114 98 6.7 6.74 Dehydration 2H2O

277�328 302 19.51 20.02 Coordination sphere 2H2O+Cl2328�342 335 10.96 11.23 Decomposition CH3CO2H(0.2L) 22.6 22.07


[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 122�223 184 8.96 8.79 Dehydration 3H2O

310�394 352 28.54 27.18 Coordination sphere 2H2O+Cl2+CH3CO2H(0.2L)



[Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 76�115 96 8.32 8.79 Dehydration 2 1/2H2O

247�301 274 27.88 27.18 Partial decomposition 2H2O+Cl2+CH3CO2H(0.19L)



[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 2 1/2 H2O 76�146 113 7.5 7.46 Dehydration 2 1/2H2O

271�400 335 17.13 17.73 Coordination sphere 2H2O+Cl2 Not complete*

400�520 460 10.95 9.95 Partial decomposition CH3CO2H(0.18L)

520�787 653 � � Final decomposition L

*T.g. curve did not reach plateau at 1000 �C.

52

mic peak at 300�C (d.t.a.); 301 and 339�C (d.s.c.). Theexothermic peaks at 483�C (d.s.c.); 557�C (d.t.a.),559�C (d.s.c.); and 564�C (d.t.a.) show final decompo-sition of the complexes.

From the d.t.a. curves, the reaction order n and Sactivation energy DEa of the thermal decompositionhave been elucidated using the Piloyan formula [38].

ln �t ¼ C� Ea=RT

The plot of ln �t versus 1000/T gave a straight linewith slope )Ea/RT (Figure 5). The obtained values aregiven in Table 5. The order of the decomposition forall complexes is unity.The activation energy values for the dehydration pro-

cesses for all complexes are more than the activation en-ergy values for the removal of lattice water [39]. Theactivation energy values for the Cu(II), Ni(II) andCo(II) complexes are expected to increase proportionallyto the decrease in their radius [40]. The activation ener-gies of the dehydration step of the complexes [Cu2 Æ -L1] Æ 2H2O, [Ni2 Æ L1] Æ 2H2O and [Co2 Æ L1] Æ 2H2Owhich have square-planar structure decrease propor-tionally to their radius degrees. DEa values for thesecomplexes are 91, 88 and 85 J/mol. The shorter the ra-dius of the metal ion, the easier the ligand approachesthe central atom. As a result, metal�ligand interactionbecomes stronger, the detachment of the link more dif-ficult, and DEa values increase [41, 42].It is shown in Table 3, that the Ni complexes are

the most stable, no matter what Schiff-base ligand ispresent. It should be also emphasized that Ni com-

Fig. 5. ln �t versus 1000/T relationship for the complex [Co2 Æ L4 ÆCl2 Æ (H2O)] Æ 2½ H2O

Table 5. D.t.a. and d.s.c. data of the complexes.

Complex Temp. range

t(�C)DTA peak

t(�C)S Ea

J/mol

Temp range

t(�C)DSC peak

t(�C)DHJ/g

Process

[Cu2 Æ L1] Æ 2H2O 107�150 135 endo 0.96 91 155�187 172 endo 157 Dehydration

200�225 213 endo 1.05 148 204�226 217 endo 32 Melting

235�251 241 exo 0.67 291 234�245 240 exo �89 Vaporization

458�700 570 exo 107 Final decomposition

½Ni2 � L1� � 2H2O 75�218 96 endo 0.85 88 � � � Dehydration

185 endo 1.03 83 217�262 239 endo 40 Dehydration


400 endo 1.08 399 397�419 400 endo 350 Partial decomposition

444�663 524 exo 78 545�577 554 exo )343 Final decomposition

[Co2 Æ L1] Æ 2H2O 98�166 135 endo 1.0 85 150�196 172 endo 142 Dehydration

276�322 293 endo 0.81 100 283�296 283 endo 75 Partial decomposition


[Ni2 Æ L2 Æ Cl2 Æ (H2O)2] Æ H2O 135�215 162 endo 1.0 191 176�207 195 endo 156 Dehydration

215�383 311 endo 1.0 46 Partial decomposition

426�629 551 exo 500�562 528 exo )343 Final decomposition

[Co2 Æ L2 Æ Cl2 Æ (H2O)2] Æ 2H2O 93�195 148 endo 1.02 73 165�215 193 endo 205 Dehydration

202�360 301 endo 1.6 248 308�323 316 endo 335 Coordination sphere

344�354 348 endo 76 Partial decomposition

478�491 483 exo )456 Final decomposition

[Ni2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 150�260 215 endo 0.81 49 215�290 258 endo 181 Dehydration


400 endo 1.0 422 386�405 396 endo 304 Partial decomposition


[Co2 Æ L3 Æ Cl2 Æ (H2O)2] Æ 3H2O 93�134 134 endo 1.0 47 127�195 174 endo 218 Dehydration



469�607 557 exo 1.2 70 550�570 559 exo )549 Final decomposition

[Co2 Æ L4 Æ Cl2 Æ (H2O)2] Æ 21/2H2O 105�178 149 endo 1.05 51 176�223 198 endo 208 Dehydration



564 exo 70 Final decomposition

53

plexes are more stable than its analogous isostructuralCu and Co complexes (Figure 5). The high stability ofNi complexes may be due to strong hydrogen bondingto the other water molecules and oxygen atom of theneighboring carboxylate group [43] which reflect thehigh dehydration temperature of the Ni complexes.Also, Ni complexes may exist as polymeric structure.In an attempt to elucidate the mechanism of the

thermal decomposition mechanism of the complexes,the i.r. spectra of the [Cu2 Æ L1] Æ 2H2O; [Ni2 Æ L1] Æ2H2O and [Co2 Æ L1] Æ 2H2O complexes heated at 150,190, 220�C; 200, 350, 400�C and 180, 300, 410�Crespectively, were recorded (Figure 6, is an example).It shows that the i.r. spectrum changes shape, intensityand position of some characteristic bands. At 150�C,the i.r. spectrum of the complex indicates that bandsat 3400, 3232 and 3082 cm)1 assigned to the �OH and�NH suffer broadening and become less intense com-pared to its i.r. spectrum at room temperature. Thiscan be explained by the formation of the hydrogenbond N�H� � �O which changes to N� � �H�O. Thismechanism can be represented as follows:

It was confirmed by heating the complex at 190�Cwhere its i.r. spectrum shows that the �NH vibrationshave completely disappeared. Also, the bands at 1587and 1429�C attributed to the presence of the carbox-ylate group are completely lost at 220�C, indicatingits decomposition and evolution as CH3COOH. Thismechanism is also applied in the case of[Ni2 Æ L1] Æ 2H2O and [Co2 Æ L1u] Æ 2H2O. Heatingcomplex [Cu2 Æ L1] Æ 2H2O induces thermochromismwith color changes which support the above-mentioned mechanism:

½Cu2 � ðC13N4O6H16Þ� � 2H2Obright green

�150 �C!

½Cu2 � ðC13N4O6H16Þ�olive green

þ 2H2O

½Cu2 � ðC13N4O6H16Þ� � 190 �C!½Cu2 � ðC11N4O4H12Þ�

brownþCH3COOH

½Cu2 � ðC11N4O4H12Þ� � 220 �C!½Cu2 � ðC11N3O4H9Þ�

deep brownþNH3

The same phenomenon was observed with [Ni2 Æ -L1] Æ 2H2O:

½Ni2 � ðC13N4O6H16Þ� � 2H2Oolive

�200 �C!

½Ni2 � ðC13N4O6H16Þ�yellow

þ 2H2O

½Ni2 � C13N4O6H16Þ� � 350 �C!½Ni2 � ðC11N3O4H9Þ�

brownþCH3COOHþNH3

½Ni2 � ðC11N3O4H9Þ� � 400 �C!½Ni2 � ðC9N3O2H5Þ�

deep brownþCH3COOH

Based on the above analytical data and physicochemicalproperties, a square planar structure was proposed forthe binuclear complexes of the L1 Schiff-base in whichthe metal ion is coordinated through the azomethinenitrogens, the amide groups and the carboxylate oxy-gen. For the binuclear complexes of L2, L3 and L4

Schiff-base, a five coordinated structure is suggested inwhich the metal atom is coordinated to the enol and car-boxylate oxygens: the azomethine and amide nitrogens.

Fig. 6. I.r. spectrum of the complex [Cu2 Æ L1] Æ 2H2O heated at dif-

ferent temperatures.

54

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schiff base complexes

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