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Spectrochimica Acta Part A 72 (2009) 421–428

Contents lists available at ScienceDirect

Spectrochimica Acta Part A: Molecular andBiomolecular Spectroscopy

journa l homepage: www.e lsev ier .com/ locate /saa

ynthesis and characterization of transition metal 2,6-pyridinedicarboxylic aciderivatives, interactions of Cu(II) and Ni(II) complexes with DNA in vitro

adaf Khana, Shahab A.A. Namia, K.S. Siddiqia,∗, Eram Husainb, Imrana Naseemb

Department of Chemistry, Aligarh Muslim University, Aligarh 202002, IndiaDepartment of Biochemistry, Aligarh Muslim University, Aligarh 202002, India

r t i c l e i n f o

rticle history:eceived 30 June 2008ccepted 10 October 2008

eywords:ononuclearipicolinic acidiperazineNA cleavagelectrophoresis

a b s t r a c t

Mononuclear complexes M(L)Cl2 where M = Mn(II), Fe(II), Co(II), Ni(II) and Cu(II) and (L = N,N-diethylpiperazinyl,2,6-pyridinedicarboxylate), have been synthesized and characterized by elementalanalysis, FT-IR, 1H NMR spectroscopy, UV–vis, magnetic moment, TGA/DSC, cyclic voltammetry and con-ductivity measurement data. The spectral data suggests that the dipicolinic acid acts as a bidentate ligandand is coordinated to the metal ion through the carboxylate oxygen. The cyclic voltammogram for Cu(L)Cl2complex was found to display two reversible Cu(II)/Cu(I) and Cu(II)/Cu(III) redox couple. The ligandexhibits a two-step thermolytic pattern while the complexes decompose in three stages respectively.An octahedral geometry has been proposed for both the complexes. The investigation of the interaction ofthe complexes with calf thymus DNA has been performed with absorption spectroscopy and fluorescence

ntibacterial activity quenching experiments, which showed that the complexes are avid binders of calf thymus DNA. Alsothe interaction of the Cu(II) and Ni(II) complexes with plasmid DNA (pUC 19) was studied using agarosegel electrophoresis. The results revealed that these complexes can act as effective DNA cleaving agentsresulting in the nicked form of DNA (pUC 19) under physiological conditions. The gel was run both in theabsence and presence of an oxidizing agent (H2O2). The ligand and its complexes have also been screenedagainst microbes in order to study their antibacterial action. The results revealed that the Cu(II) complex

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has activity comparable w

. Introduction

Pyridine-2,6-dicarboxylic acid or dipicolinic acid (DPA) is knowno form stable chelates with metal ions and oxometal cations andan display widely varying coordination demeanour functioning asbidentate, tridentate, meridian or bridging ligand [1,2]. DPA can

tabilize unusual oxidation states and can form bridging hydrogenonds due to its functional groups [3]. It is a versatile N,O-chelatinggent with limited stearic hindrance and can further provide possi-ility to form polymeric complexes through bridging coordinationf carboxylates under suitable conditions [4]. DPA can act as annteresting ligand due to its ability to form strong covalent bonds.lso the spatial separation of two carboxylate groups attached

o the same aromatic ring leads to either a polymeric chain or a

yclo-oligomeric ring structure and there is a potential influence ofitrogen atom on the coordination mode [5].

A very important characteristic of these ligands is their diverseiological activity [6]. They are present in many natural products,

∗ Corresponding author.E-mail address: kssiddiqi@yahoo.co.in (K.S. Siddiqi).

atictsCc

386-1425/$ – see front matter © 2008 Elsevier B.V. All rights reserved.oi:10.1016/j.saa.2008.10.001

e reference drugs gentamycin and flucanzole.© 2008 Elsevier B.V. All rights reserved.

s an oxidative degradation product of vitamins, coenzymes andlkaloids and are also an important constituent of fulvic acids. DPAomplexes of iron are well-known electron carriers in various bio-ogical models and are recognized as specific molecular tools forNA cleavage [7]. DPA is also an intermediate in the tryptophanegradation pathway and is a precursor for NAD [8].

Transition metal complexes have been the subject of thoroughnvestigation because of their extensive applications in wide rang-ng areas from material sciences to biological sciences [9]. Metalomplexes are well-known to accelerate the drug action and thefficacy of a therapeutic agent can often be enhanced upon coor-ination with a metal ion [10]. The pharmacological activity haslso been found to be highly dependent on the nature of theetal ion and the donor sequence of the ligands as different lig-

nds exhibit different biological properties [11]. In recent years,he binding studies of transition metal complexes have become anmportant field in the development of DNA molecular probes and

hemotherapeutics [12]. There is a substantial literature supportinghe application of artificial DNA cleaving agents in biotechnology;tructural studies of nucleic acids or development of new drugs.ompounds showing the properties of effective binding as well asleaving double stranded DNA under physiological conditions are

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f much importance since these could be used as diagnostic agentsn medicinal and genomic research. With the recognition of cis-latin as the anticancer drug the introduction of metal-based drugsas gained importance. However, the greatest disadvantage of cis-latin and other heavy metal-based drugs is their toxicity [13]. Itill be more appropriate to use biologically soft metal ions and

heir complexes, reducing the toxic effect and enabling faster andfficient removal of the drug from the body. Transition metal com-lexes interact with DNA through covalent binding, electrostatic

nteractions, groove binding or intercalation [14,15]. Studies haveroved that macrocyclic complexes as chemical nucleases are thefficient DNA cleavage agents. Most of the enzymes that partici-ate in the biochemistry of nucleic acids require divalent metal ionofactors to promote their activity for the cleavage of phosphodi-ster bond of DNA [16].

Metal complexes of pyridine carboxylic acids and some of theirerivatives have also been used as model systems for the designf new metallopharmaceutical compounds [17]. Lately DPA hasecome one of the more suitable ligands for modelling potentialharmacologically active compounds because of its low toxicity andmphophilic nature [18].

Copper being a bio-essential element its complexes find morepplications in nucleic acid chemistry as compared to other heav-er transition metal elements [19]. Over a dozen of enzymes thatepend on copper, for their activity have been identified; theetabolic conversions catalyzed by all these enzymes are oxida-

ive in nature [20]. Due to their importance in biological processes,u(II) complexes synthesis and activity studies have been the

ocus from different perspectives. Recently, the studies have provedhat various macrocycles of copper as chemical nucleases are effi-ient DNA cleaving agents [21]. Due to the versatile biological andhemical properties of DPA we are reporting the synthesis andharacterization of a ligand L where L = N,N-diethylpiperazinyl,2,6-yridinedicarboxylate and its mononuclear complexes with Mn(II),e(II), Co(II), Ni(II) and Cu(II) ions. Since Cu(II) and Ni(II) are bio-hemically active their interaction with pUC 19 plasmid DNA haseen investigated in order to ascertain the probable mechanism ofheir cleavage under physiological conditions.

. Experimental

.1. Chemicals and methods

Hydrated metal chlorides, 2,6-pyridinedicarboxylic acidAldrich), dichloroethane (Merck), piperazine hexahydrate (Lobahemie) were used as received. CT-DNA and pUC 19 plasmid

NA were purchased from Sigma. Methanol was distilled prior tose. Elemental analyses (C, H, N and S) were carried out with aarlo Erba EA-1108 analyzer. The metal contents were estimatedy complexometric titration [22]. IR spectra (4000–400 cm−1)ere recorded on a RXI FT-IR spectrometer as KBr disc while

flqlat

Scheme 1. Schematic diagram for t

Part A 72 (2009) 421–428

he 600–200 cm−1 range was scanned with CsI on a Nexus FT-IRhermo Nicolet (Wisconsin). The Electronic spectra were recordedn a Cintra 5GBC spectrophotometer in DMSO. The NMR spectraere recorded on a DPX-300 spectrometer in DMSO at room

emperature. The conductivity measurements were carried outn a CM-82T Elico conductivity bridge in DMSO. Magnetic sus-eptibility measurements were done with a 155 Allied Researchibrating sample magnetometer at room temperature. TGA waserformed with a PerkinElmer (Pyris Diamond) thermal analyzernder nitrogen atmosphere using alumina powder as reference.he weight of the sample was kept between 8 and 12 mg and theeating rate was maintained at 10 ◦C/min. Cyclic voltammetryas performed on a 0.1 mM DMSO solution. The potential was

canned at variable rates from 25–100 mV/s. The potentiometeras a Princeton Applied Research model 263A attached to a

.C. using Electrochemistry Powersuite Software. The workinglectrode was a glassy carbon electrode, the counter electrodeas a platinum wire and the reference electrode consisted of Ag,gCl/KCl (saturated).

.2. Synthesis

.2.1. Synthesis of the ligand LL = N,N-diethylpiperazinyl,2,6-pyridinedicarboxylate)

DPA (2 mmol, 0.33 g) was dissolved in 15 mL of methanol.ichloroethane (4 mmol, 0.35 mL) was dissolved in 5 mL of the

ame solvent. The dissolved dichloromethane was then carefullydded drop by drop into the methanolic solution of DPA under con-inuous stirring. This reaction mixture was then refluxed on a hotlate for 3 h. A white precipitate appears at this time which dis-olves and becomes a clear solution on further refluxing of 4 h. Thislear solution was then brought to room temperature. Piperazineexahydrate (2 mmol, 0.38 g) was dissolved in 15 mL methanol andhen it was added to the clear solution form of reaction mixturelowly accompanied by continuous stirring on a magnetic stirrer.he addition of piperazine hexahydrate gives an immediate pre-ipitation. The reaction mixture in the form of a white precipitateas further refluxed for 4 h in order to ascertain complete precip-

tation. Finally after refluxing, the white precipitate obtained wasltered using a Whatman filter paper (no. 100). It was then washedith methanol and dried in vacuo (Scheme 1). The progress of the

eaction was monitored by thin layer chromatography (TLC).

.2.2. Synthesis of complexes

.2.2.1. Synthesis of the complex Mn(L)Cl2 where L = N,N-iethylpiperazinyl,2,6-pyridine dicarboxylate. The ligand (2 mmol,.60 g) was dissolved in 15 mL of hot methanol in a round bottom

ask. Metal halide (2 mmol, 0.39 g) was dissolved in minimumuantity of the same solvent. It was then gradually added to theigand solution with stirring, to obtain an immediate precipitationccompanied by a visible colour change. The reaction mixture washen refluxed for 6 h on a hot plate, which affords the precipitate.

he preparation of the ligand.

S. Khan et al. / Spectrochimica Acta Part A 72 (2009) 421–428 423

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Scheme 2. Schematic diagram for the preparation of the m

he precipitation of the complex ensures the completion of theeaction. The complex was then carefully filtered by using What-an filter paper (no. 100), washed with methanol and dried in

acuo (Scheme 2).

.2.2.2. Synthesis of the complex Fe(L)Cl2 where L = N,N-iethylpiperazinyl,2,6-pyridinedicarboxylate.. The ligand (2 mmol,.60 g) was dissolved in 15 mL of hot methanol in a round bottomask. Metal halide (2 mmol, 0.39 g) was dissolved in minimumuantity of the same solvent. It was then gradually added to theigand solution with stirring, to obtain an immediate precipitationccompanied by a visible colour change. The reaction mixture washen refluxed for 6 h on a hot plate, which affords the precipitate.he precipitation of the complex ensures the completion of theeaction. The complex was then carefully filtered by using What-an filter paper (no. 100), washed with methanol and dried in

acuo (Scheme 2).

.2.2.3. Synthesis of the complex Co(L)Cl2 where L = N,N-iethylpiperazinyl,2,6-pyridinedicarboxylate.. The ligand (2 mmol,.60 g) was dissolved in 15 mL of hot methanol in a round bottomask. Metal halide (2 mmol, 0.47 g) was dissolved in minimumuantity of the same solvent. It was then gradually added to theigand solution with stirring, to obtain an immediate precipitationccompanied by a visible colour change. The reaction mixture washen refluxed for 6 h on a hot plate, which affords the precipitate.he precipitation of the complex ensures the completion of theeaction. The complex was then carefully filtered by using What-an filter paper (no. 100), washed with methanol and dried in

acuo (Scheme 2).

.2.2.4. Synthesis of the complex Ni(L)Cl2 where L= N,N-iethylpiperazinyl,2,6-pyridinedicarboxylate.. The ligand (2 mmol,.60 g) was dissolved in 15 mL of hot methanol in a round bottomask. Metal halide (2 mmol, 0.47 g) was dissolved in minimumuantity of the same solvent. It was then gradually added to theigand solution with stirring, to obtain an immediate precipitationccompanied by a visible colour change. The reaction mixture washen refluxed for 6 h on a hot plate, which affords the precipitate.he precipitation of the complex ensures the completion of theeaction. The complex was then carefully filtered by using What-an filter paper (no. 100), washed with methanol and dried in

acuo (Scheme 2).

.2.2.5. Synthesis of the complex Cu(L)Cl2 where L = N,N-iethylpiperazinyl,2,6-pyridinedicarboxylate.. The ligand (2 mmol,.60 g) was dissolved in 15 mL of hot methanol in a round bottomask. Metal halide (2 mmol, 0.34 g) was dissolved in minimumuantity of the same solvent. It was then gradually added to the

1nwb1

omplexes where M = Mn(II), Fe(II), Co(II), Ni(II), and Cu(II).

igand solution with stirring, to obtain an immediate precipitationccompanied by a visible colour change. The reaction mixture washen refluxed for 6 h on a hot plate, which affords the precipitate.he precipitation of the complex ensures the completion of theeaction. The complex was then carefully filtered by using What-an filter paper (no. 100), washed with methanol and dried in

acuo (Scheme 2).

.3. Interaction of the complexes with DNA–fluorescencepectroscopy

The DNA binding studies were carried out in 0.02 mol/L of phos-hate buffer containing 60 mM NaCl at pH 7.0 at room temperature.alf thymus DNA as sodium salt (CT-DNA) was purchased fromigma and the purity was checked by the absorbance A260/A280hich was found to be 1.86 in phosphate buffer [23] at pH 7.0.

he concentration of ethidium bromide (EB), which binds to DNAy interacting and enhancing the fluorescence intensity, was keptt 4 �M. The concentration of DNA (25 �M) and EB was kept con-tant. The fluorescence studies were carried out with increasingoncentration of the complexes ranging from 5 to 25 �M. Since theetal complexes do not fluoresce themselves a fluorophore like EBas utilized. All the samples were excited at 340 nm. Emission was

ecorded at 500–600 nm. The emission slit was 10 nm.

.4. Interaction of the complexes with CT-DNA-absorptionpectroscopy

DNA stock solutions (5 mM) were prepared by dilution of CT-NA with buffer (containing 150 mM NaCl and 15 mM trisodiumitrate at pH 7.0) followed by exhaustive stirring at 4 ◦C for 3 days24] and the solutions were kept at 4 ◦C for a week. The nucleotideoncentrations were determined by the absorption at 260 nm using= 6600 M−1 cm−1 (expressed as phosphate).

.5. Assay of nuclease activity

The DNA cleavage by the complexes was probed using pUC 19upercoiled plasmid DNA employing agarose gel electrophoresis.he metal complexes were dissolved in 50 mM Tris–HCl/18 mMaCl buffer (pH 7.2). Increasing concentrations of the metal com-lexes 1 �L from different concentrations (5–25 �M) were taken inlean eppendorf tubes and plasmid DNA (1 �L of 0.10 �g/mL) wasdded. The contents were incubated at 37 ◦C for 12 h and loaded on

% agarose gel after mixing 5 �L of loading buffer (25% bromophe-ol, 30% glycerol (3 �L) and 0.25% xylene cyanol). Electrophoresisas performed at a constant voltage (80 V) till the bromophenollue reached to three-fourth of the gel. The gel was stained for0 min. by immersing it in ethidium bromide solution. It was then

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24 S. Khan et al. / Spectrochimi

e-stained by keeping in sterile distilled water for 10 min. Plasmidand was visualized by viewing the gel under transilluminator andhotographed. The efficiency of the DNA cleavage was measuredy determining the ability of the complex to form open circular oricked forms, respectively.

.6. Pharmacology

The antibacterial and antifungal activity of a solution of freshlyynthesized complexes in dichloromethane was tested againstome gram positive and gram negative bacteria such as, E. coli, S.ureus, C. albicans and A. flavus. The activity was then comparedith some reference antibiotics that were purchased from the mar-

et. All the complexes and the parent ligand were screened for theirctivity against the test organisms. The hole plate diffusion methodas adopted for the activity measurements. The bacterial strainsere grown in nutrient agar slants and the fungal strains were

rown in Sabouraud dextrose agar slants. The viable bacterial cellsere swabbed onto Nutrient agar plates and the fungal spores onto

abouraud dextrose agar plates. The compounds were dissolved inichloromethane to a final concentration of 0.1%. A 0.5 cm diameterell was cut in a medium inoculated with the respective cultures,

nd the test solutions of the compounds in different concentrations5 and 1 �g) for bacterial cultures and (15 and 50 �g) for fungalultures were allowed to stay in the wells. The petri plates werencubated for 36 h for bacterial cultures and 72 h for fungal cultures.ll the compounds were screened against flucanazole as reference

or fungal cultures and gentamycin for bacterial cultures in theirtandard concentration of 200 �g/well. The activity of the com-ounds was evaluated by measuring the diameter of the inhibitionone around the respective wells.

. Results and discussion

The complexes were obtained in moderate yield by the reactionf metal chloride, MX2.nH2O with ligand, L in methanolic mediaScheme 2).

X2·nH2O + L → M(L)X2

here M = Mn(II), Fe(II), Co(II), Ni(II), Cu(II), X = Cl and L = N,N-iethylpiperazinyl,2,6-pyridinedicarboxylate. Elemental analysis,T-IR, UV–vis spectroscopy, 1H NMR, cyclic voltammetry andGA/DSC were used to characterize the complexes. The com-lexes are soluble in H2O and common organic solvents. Theonductivity measurement (10−3 M) in DMSO indicated them toe non-electrolytes [25] (Table 1).

The interaction of Ni(L)Cl2 and Cu(L)Cl2 complexes with pUC9 was studied by monitoring the conversion of circular super-

oiled DNA (Form I) to nicked DNA (Form II). The amount of strandcission was assessed by agarose gel electrophoresis. The interca-ating ability of the complexes was probed by DNA–EB system usinguorescence quenching spectroscopic experiments which suggestshat the Cu(II) complex is a better intercalator than the Ni(L)Cl2.

Mbtb1

able 1hysicochemical properties of the ligand and its complexes.

ompounds Colour Yield (%) Molar conductance (�−1 mol−1 cm2) Anal

C

White 80 17 58.7n(L)Cl2 Pale yellow 62 27 41.9

e(L)Cl2 Brown 67 42 41.4o(L)Cl2 Grey 70 17 41.6i(L)Cl2 Green 68 11 41.4u(L)Cl2 Blue 55 29 40.8

Part A 72 (2009) 421–428

.1. Infrared spectroscopy

Various coordinating modes exists in the pyridine carboxyliccid complexes which can be distinguished from the differencen asymmetric and symmetric carboxylate [26] stretching bands,

� = �asCOO− −�symCOO−. If there is a marked difference of over00 cm−1 it implies monodentate coordination while a differencef about 100 cm−1 confirms a bidentate coordination. The carboxy-ate ion may coordinate to a metal ion in one of the following ways.

In the IR spectra of the complexes the �(C–O) bands decrease inntensity significantly as compared to free ligand, L and new bandsppear at 1623, 1625 and 1415, 1412 cm−1 corresponding to �as ands frequencies, respectively. The large difference between �as ands (�� > 200 cm−1) frequencies indicates monodentate coordina-ion [27]. The absence of any band in the 3000–3450 cm−1 regionmplies deprotonation of carboxylate groups leading to a ring liketructure (Scheme 1) [28]. Furthermore, the absence of any band in000–3450 cm−1 region can also be corroborated with the depro-onation of piperazine. Piperazine exists in two conformations i.e.hair and boat form although the chair form has a higher stability.owever, the boat form is predominant [29] in case of bidentate

helation with the same metal ion such as [PtCl2(Me2ppz)]. In ourase too, the piperazinato moiety is constrained to form a ring liketructure thus existing in a boat conformation (Table 2) [30].

.2. NMR spectroscopy

The 1H NMR spectrum of the free DPA in DMSO exhibits a sharpinglet at ı 10.14 ppm due to the carboxylic acid protons, which isound to be absent in metal complexes indicating the formation of–O–C bonds [31]. The eight piperazinium protons were found toesonate at 3.15 ppm as a singlet in case of ligand and do not alterfter complexation [32]. The methylene protons were observed at.64 ppm in case of ligand and are shifted slightly when complexedith metal ions (2.66–2.69 ppm) [33].

.3. Electronic spectroscopy and magnetic moment

The electronic spectral bands and the magnetic moments ofhe complexes are listed in Table 3. In octahedral environment,

n(II) complex gives spin forbidden as well as parity-forbiddenands. In addition to the band due to n–�* transition, the elec-ronic spectrum of Mn(II) complex in DMSO exhibits three moreands in the region 31104–30350 cm−1; 22542–20533 cm−1 and7513–16890 cm−1 which have been assigned to 4T1g(P)← 6A1g;

ysis, (%) Found (calculated)

H N O M

8 (59.0) 6.19 (6.27) 13.79 (13.76) 20.74 (20.95) –3 (41.74) 4.53 (4.43) 9.48 (9.74) 14.55 (14.84) 12.82 (12.74)8 (41.69) 4.57 (4.43) 9.66 (9.72) 15.08 (14.81) 12.85 (12.92)3 (41.40) 4.28 (4.40) 9.43 (9.65) 14.81 (14.70) 13.48 (13.54)6 (41.42) 4.35 (4.40) 9.44 (9.66) 14.86 (14.71) 13.68 (13.49)7 (40.96) 4.22 (4.35) 9.41 (9.55) 14.75 (14.55) 14.67 (14.44)

S. Khan et al. / Spectrochimica Acta Part A 72 (2009) 421–428 425

Table 2Characteristic IR bands (in cm−1) of the ligand and its complexes.

Complex �asym �sym Ring vibration �(C–N) �(M–X)

L 1724 – 1300 w, 1270 m 1634 –Mn(L)Cl2 1623 1415 1288 w, 1262 m 1598 380 wFe(L)Cl2 1625 1412 1285 w, 1264 m 1606 367 wCo(L)Cl2 1629 1425 1291 w, 1266 m 1604 355 wNi(L)Cl2 1631 1431 1284 w, 1254 m 1588 340 wCu(L)Cl2 1635 1418 1278 w, 1244 m 1586 327 w

Table 3Magnetic susceptibility, electronic spectra and ligand field parameters of the complexes.

Complex Magnetic moment (B.M.) Electronic bands (cm−1) Log ε (l mol−1 cm1) Possible assignments 10 Dq (cm−1) B (cm−1) ˇ

Mn(L)Cl2 5.77 22430 3.2 4A1(G)← 6A1 20075 944 0.7520640 2.7 4T1(G)← 6A1

Fe(L)Cl2 5.32 23529 4.1 5E← 5T2 – – –Co(L)Cl2 4.52 17460 3.0 4T1(P)← 4A2(F) – – –

11380 2.1 4T1(F)← 4A2(F)N 1 1

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tirgCtmslpr0respectively (E1/2 = (Epa + Epc)/2). The criterion of reversibility wasverified by the ratio, ipa/ipc = 0.89–0.92 �A with variation of scanrates. Any redox behaviour of the ligand in the range−1.25 to 1.50 Vis ruled out because dipicolinic acid shows quasi-reversible andirreversible process at −1.58 V [45].

i(L)Cl2 Diamagnetic 23696 3.215503 2.8

u(L)Cl2 2.13 18529 1.814814 1.2

T2g(G)← 6A1g and 4T1g(G)← 6A1g transitions, respectively. Theigh spin d5 configuration gives an essentially spin-only mag-etic moment of ∼5.9 B.M., which is temperature independent.he Mn(II) complex under consideration has the value of 5.84.M. which is close to the calculated value. Thus the ligand fieldands and magnetic moment value support an octahedral geometryround the metal ion [34].

For the octahedral spin-free Fe(II) complexes the magneticoment values lies at about 5.5 B.M. and is nearly independent

f temperature. In the present case the magnetic moment valuesre found to be 5.30 B.M. for Fe(L)Cl2 complex. The deviation in theagnetic moment value may be ascribed to a slight distortion from

he regular octahedral geometry [35] (Table 3).Three intense bands exhibited by the Co(L)Cl2 at 15480; 20781

nd 24272 cm−1 are attributed to the three transitions from 4T1g(F)round state to 4T2g(F), 4A2g(F) and 4T1g(P) states, respectivelyhich is characteristic of the high-spin octahedral Co(II) complexes

36]. Dq = 1580 cm−1 and B = 753 cm−1 are also in accordance withhe octahedral environment of Co(II) complexes [37]. The roomemperature magnetic moment of 4.39 B.M. further supports anctahedral arrangement around Co(II) ion in Co(L)Cl2.

Generally, bivalent octahedral Ni(II) complexes exhibit threepin allowed d–d bands which shift to a higher wave num-er depending upon the nature of the coordinating ligand [38].he electronic spectrum of the Ni(L)Cl2 complex in aqueousedium exhibits two absorption bands in the region 9560–9090

nd 16722–16129 cm−1 corresponding to the 4T2g← 4A2g andT1g(F)← 4A2g transitions, respectively, which is characteristic ofn octahedral Ni(II) ion [39]. A broad hump observed at 12500 cm−1

s also present and is probably due to the spin forbidden transi-ion similar to that observed by Prushan et al. for the Ni(L)Cl2 andi(L)(NO3)2 where L = trithiatridecane-2,12–dionedioxime [40].enerally the magnetic moment for octahedral Ni(II) complex, liesetween 2.9 and 3.4 B.M. Our value (3.26 B.M) is well within thepecified limit therefore, an octahedral geometry is proposed forhe Ni(II) ion.

The blue colored Cu(L)Cl2 complex in the present case exhibits

broad band centered at 730 nm encompassing several tran-

itions characteristic of an octahedral geometry [41]. However,everal attempts to resolve this band were unsuccessful due tohe absence of cubic field symmetry. The magnetic moment of

ononuclear Cu(II) complexes generally lies between 1.75 and 2.20

B1g← A1g 15560 1613 0.641A2g← 1A1g2A1g← 2B1g – – –2Eg← 2B1g

.M. regardless of the stereochemistry and is temperature indepen-ent. However, the room temperature magnetic moment of 2.11.M. in our case is in agreement with that found for octahedralu(RPO)py2 complex [42].

.4. Cyclic voltammetry

The complexes are soluble in common organic solvents andhe electron transfer behavior of the Cu-complex was studiedn CH3CN by two-vertex cyclic voltammetry in −2.0 to +2.0 Vange with variable scan rate (25–100 mV/s). The cyclic voltammo-ram at Pt disk electrode for this complex display two reversibleu(II)/Cu(I) and Cu(II)/Cu(III) redox couple (Fig. 1). It is knownhat the redox potential of Cu(II)/Cu(I) process is shifted towards

ore negative potential as the electron-donating ability of theubstituents on the ligand framework becomes higher [43]. Theigand is electrochemically active due to the presence of theyridinium moiety [44]. The copper(II) ion yields two quasi-eversible redox couple, located at E1

1/2 = 0.445 V and E21/2 =

.805 V which are assignable to Cu(II)/Cu(I) and Cu(II)/Cu(III),

Fig. 1. Cyclic voltammogram of Cu(L)Cl2.

426 S. Khan et al. / Spectrochimica Acta Part A 72 (2009) 421–428

Table 4Thermal degradation of ligand and its complexes.

Complex First decomposition stage Second decomposition stage Third decomposition stage Residue

Fragments Temperaturerange (◦C)

Mass lossdata, found(calculated)(%)

Fragments Temperaturerange (◦C)

Mass lossdata, found(calculated)(%)

Fragments Temperaturerange (◦C)

Mass lossdata, found(calculated)(%)

L C8H16N2 162–255 46.31 (45.93) C7H3NO4 260–400 52.44 (54.07) – – – –Mn(L)Cl2 2HCl 182–220 16.91 (16.76) C8H14N2 225–330 31.11 (31.77) C7H3NO3 360–470 33.87 (34.28) MnOFe(L)Cl2 2HCl 180–224 17.02 (16.58) C8H14N2 225–310 31.25 (31.43) C7H3NO3 360–470 33.57 (33.90) FeOC 30 31.11 (31.77) C7H3NO3 360–470 33.87 (34.28) CoON 10 31.25 (31.43) C7H3NO3 360–470 33.57 (33.90) NiOC 30 31.11 (31.77) C7H3NO3 360–470 33.87 (34.28) CuO

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wiDftDbtmtsiatNs

F(EE

pireplaced by the Cu(II) complex (Fig. 2). A similar treatment withNi(II) complex showed a less pronounced effect indicating thatthe Cu(II) complex is a better intercalator than the Ni(II) analogue(Fig. 3).

o(L)Cl2 2HCl 182–220 16.91 (16.76) C8H14N2 225–3i(L)Cl2 2HCl 180–224 17.02 (16.58) C8H14N2 225–3u(L)Cl2 2HCl 182–220 16.91 (16.76) C8H14N2 225–3

.5. Thermal studies (TGA/DSC)

The thermal analysis of the ligand and its mononuclear com-lexes was carried out under an inert atmosphere (N2). The

igand exhibits a two-step thermolytic pattern while the com-lexes decompose in three stages, respectively (Table 4). Theecomposition of the free ligand begins at 162 ◦C and exhibits aharp endotherm (absent in the complexes) at 195 ◦C. However, inhe case of complexes the first decomposition spans from 180 ◦Chrough 220 ◦C corresponding to the liberation of chloride ion asCl (found ∼17%). The results show a significant increase in theecomposition temperature when compared with the ligand, indi-ating a higher thermal stability of the complexes [46]. The absencef any thermal change before 150 ◦C indicates that the restructuringf the ligand and its complexes does not take place before the ini-iation of the degradation process and also rules out the possibilityf any water or solvent molecule in the ligand and its complexes47]. The second thermolytic step starts at 220 ◦C and terminatest about 360 ◦C, corresponding to the decomposition of piperazinend ethane linkage. The third stage of pyrolysis corresponds to4% weight loss consistent with the expulsion of dipicolinic moiety

eaving behind metal oxide as the end product [48].The DSC plots are consistent with the decomposition pattern of

he ligand and its complexes. Two broad peaks observed at 195 and65 ◦C indicate the pyrolysis of the ligand as an exothermic process.sharp exothermic peak obtained for the liberation of chloride

eflects fast decomposition process, implying volatilization of theomplex upon heating. However there is no well-defined peak forhe formation of metal oxide as the end product.

.6. Interaction of the complexes with Calf thymusNA-fluorescence spectroscopy

EB is one of the most sensitive fluorescent probes that can bindith DNA. There is an increase in the fluorescence of EB due to

ntercalation with DNA. If the metal complex intercalates withNA it leads to a decrease in the binding sites of DNA available

or EB, which is reflected, from a decrease in the fluorescence ofhe EB-DNA system [49]. Generally metal complexes interact withNA via intercalation or through covalent binding of the nitrogenases. However, due the presence of benzene ring, an intercala-ive binding mode may be more plausible. The metal complex

ight have penetrated between the double helical structure of DNAhrough hydrogen bonding leading to a decrease in the bindingites of DNA–EB system. Consequently a decrease in fluorescence

s observed on addition of metal complex in the DNA–EB systemnd the effect is enhanced with the increase in the concentra-ion of the complex. Accordingly, the titration of DNA–EB withi(II) and Cu(II) complexes decreased the intensity of fluorescence

pectra in our system. The addition of Cu(L)Cl2 complex to DNA

FCEE

ig. 2. Emission spectra of Cu(L)Cl2 complex (excited at 550 nm, 5 ml solution):a) control (4 �M EB + 25 �DNA); (b) 4 �M EB + 25 �DNA + 5 �M Cu(L)Cl2; (c) 4 �MB + 25 �DNA + 10 �M Cu(L)Cl2; (d) 4 �M EB + 25 �DNA + 15 �M Cu(L)Cl2; (e) 4 �MB + 25 �DNA + 20 �M Cu(L)Cl2; (f) 4 �M EB + 25 �DNA + 25 �M Cu(L)Cl2.

retreated with EB causes appreciable reduction in the emissionntensity indicating that the DNA bound EB fluorophore is partially

ig. 3. Emission spectra of Ni(L)Cl2 complex (excited at 550 nm, 5 ml solution): (a)ontrol (4 �M EB + 25 �DNA); (b) 4 �M EB + 25 �DNA + 10 �M Ni(L)Cl2; (c) 4 �MB + 25 �DNA + 15 �M Ni(L)Cl2; (d) 4 �M EB + 25 �DNA + 20 �M Ni(L)Cl2; (e) 4 �MB + 25 �DNA + 25 �M Ni(L)Cl2.

S. Khan et al. / Spectrochimica Acta Part A 72 (2009) 421–428 427

Fig. 4. Agarose gel electrophoresis diagram showing the cleavage of pUC 19 plasmidDt(D

3

umiftwccfficsicotbaTt

cHmtbH

C

dpcd

Fig. 5. Agarose gel electrophoresis diagram showing the cleavage of pUC 19 plasmidDNA by Cu(L)Cl2 in various concentrations in TAE buffer at pH 7.2. lane (a) DNA con-t(lD

ittr

3s

CCocaptot(

3.9. Antibacterial activity

The results of the antimicrobial activity are presented in theTable 5. The complexes display stable inhibition zones and are more

NA by Ni(L)Cl2 in various concentrations in TAE buffer at pH 7.2. lane (a) DNA con-rol showing pUC 19 plasmid DNA in supercoiled form, form I; lane (b) DNA + Ni(L)Cl25 �M); lane (c) DNA + Ni(L)Cl2 (10 �M); lane (d) DNA + Ni(L)Cl2 (15 �M); lane (e)NA + Ni(L)Cl2 (20 �M) lane (f) DNA + Cu(L)Cl2 (20 �M).

.7. Interaction of complexes with pUC 19 plasmid DNA

The DNA cleavage efficiency of metal complexes was examinedsing supercoiled pUC 19 plasmid DNA as the target. Circular plas-id DNA is ideally suited to probe the cleavage as the DNA exists

n a supercoiled state in its native form and converts to a relaxedorm upon single strand scission. This is exhibited by altered migra-ion rate during agarose gel electrophoresis, the fastest migrationill be observed for the supercoiled form (Form I). If one strand is

leaved, the supercoils will relax to produce a slower moving openircular form (Form II) [50]. If both the strands are cleaved a linearorm will be generated that migrates in between. In the quest tond out whether the Ni(II) and Cu(II) complexes would exhibit DNAleavage activity in vitro their effects at various concentrations weretudied using pUC 19 plasmid DNA and the results are summarizedn Figs. 4 and 5. It was found that both the Ni(II) and Cu(II) complexesan effectively cleave the supercoiled DNA, but the nuclease activityf the later complex was pronounced. This behavior is verified fromhe gel electrophoresis experiments (Fig. 4) where copper shows aetter nuclease activity as compared to nickel at the highest equiv-lent concentration (20 �M) leading to a pronounced DNA damage.he greater nuclease activity of Cu(II) complex may be attributedo its redox nature i.e. the ease of reduction of Cu(II) to Cu(I).

It was also observed that the DNA cleaving ability of Cu(L)Cl2omplex was further enhanced in the presence of an oxidant like2O2. It is evident from Fig. 5 that the copper complex cleaves DNAore efficiently in the presence of an oxidant, which is attributed to

he formation of hydroxyl radicals. The hydroxyl radical may haveeen produced due to the reaction between the metal complex and2O2 as shown below

u+ +H2O2→ Cu2+ +OH + OH•

•

These OH free radicals participate in the oxidation of theeoxyribose moiety, followed by hydrolytic cleavage of the sugarhosphate backbone. The more pronounced nuclease activity inase of Cu(L)Cl2 complex may be due to the greater ease of its oxi-ation. The increase in concentration in each well led to a decrease

FDb

rol showing pUC 19 plasmid DNA in supercoiled form, form I; lane (b) DNA + Cu(L)Cl25 �M); lane (c) DNA + Cu(L)Cl2 (5 �M) + H2O2; lane (d) DNA + Cu(L)Cl2 (10 �M);ane (e) DNA + Cu(L)Cl2 (10 �M) + H2O2 lane (f) DNA + Cu(L)Cl2 (20 �M); lane (g)NA + Cu(L)Cl2 (20 �M) + H2O2.

n the intensity of supercoiled DNA band although the intensity ofhe nicked DNA increases apparently. These observations confirmhat the Cu(II) complex acts as a better nuclease at a physiologicallyelevant concentration.

.8. Study of the complexes with CT-DNA-absorptionpectroscopy

The absorption spectra of the interaction of the Ni(L)Cl2 andu(L)Cl2 complexes with CT-DNA have been recorded for a constantT-DNA concentration of 3.125×10−4 M in presence and absencef an oxidant, H2O2. Representative spectra of Ni(L)Cl2 and Cu(L)Cl2omplexes having a concentration of 4×10−6 M along with CT-DNAlone is given in Fig. 6. The spectra clearly shows that Cu(L)Cl2 com-lex is a better intercalator than its Ni analogue which might be dueo the redox nature of Cu(L)Cl2 complex. The DNA cleavage activityf Cu(L)Cl2 complex is further enhanced in presence of H2O2 dueo the increase in hydroxyl ion concentration by Fenton reactionFig. 6).

ig. 6. Absorption spectra of Ni(L)Cl2 and Cu(L)Cl2 complexes in the presence of CT-NA. The r value is 1:100. The spectra were recorded at 25 ◦C after complexes hadeen incubated with CT- DNA for 24 h at 37 ◦C.

428 S. Khan et al. / Spectrochimica Acta Part A 72 (2009) 421–428

Table 5Inhibition zone in mm.

Compounds E. coli S. aureus A. flavus C. albicans

5 �g 1 �g 5 �g 1 �g 5 �g 1 �g 5 �g 1 �g

L 14.2 11.2 15.7 13.2 15.4 13.7 10.4 10.2Mn(L)Cl2 15.6 13.6 21.5 20.3 11.0 10.0 16.9 14.8Fe(L)Cl2 11.2 10.2 17.5 15.8 9.5 9.0 10.2 10.0CNCG

aaftetcipwatcfntaaTtcb

R

[

[

[[[

[[[

[

[[

[

[[

[

[[[

[

[

[

[[

[

[

[

[[[

[

[[

[[

[[

[

[

[[

o(L)Cl2 21.9 20.2 21.9i(L)Cl2 23.5 20.3 27.3u(L)Cl2 25.0 22.6 28.7entamycin/flucanazole 26.0 26.0 30.0

ctive than the ligand. They possess potential inhibitory activity inmounts as low as 1 �g/well for bacterial cultures and 15 �g/wellor fungal cultures. The activity of the complexes is measured inerms of inhibition of the replication of DNA by interacting with thenzyme prosthetic group. The antibacterial results evidently showhat the activity of the ligand became more pronounced and signifi-ant when coordinated to the metal ion (Table 5). This enhancementn the activity may be due to an efficient diffusion of the metal com-lexes in to the bacterial cell or interaction with the bacterial cellalls. Also it has been observed that the activity of Mn(II), Fe(II)

nd Co(II) complexes is lower than that of Cu(II) complex, whilehat of Cu(II) and Ni(II) analogues is quite comparable. The antimi-robial activity of the metal complexes generally depends on theollowing factors: the chelation ability of the ligand, the nature ofitrogen donor ligands, the total charge of the complex, the exis-ence and the nature of the metal ion neutralizing the ionic complexnd the nuclearity of the metal center in the complex [51]. Higherctivity of Cu(II) complex is probably due to the effective chelation.his activity is quite comparable with the reference drugs gen-amycin/flucanzole. However, the reduced activities in some casesan be attributed to the inability of the complexes to form hydrogenonds with the cell constituents [52].

eferences

[1] M.C. Grossel, C.A. Golden, J.R. Gomn, P.N. Horton, D.A.S. Merckel, M.E. Oszer,R.A. Parker, Cryst. Eng. Commun. 42 (2001) 1–4.

[2] Y. Ducommun, L. Helm, G. Laurenezy, A. Merbach, Inorg. Chim. Acta 158 (1989)3–4.

[3] D.L. Hoff, D.G. Tisley, R.A. Walton, J. Chem. Soc. Dalton Trans. (1973) 200–204.[4] A. Szorcsik, L. Nagy, J. Sletten, G. Szalontai, E. Kamu, T. Fiore, L. Pellerito, E.

Klman, J. Organomet. Chem. 689 (2004) 1145–1154.[5] C. Ma, J. Li, R. Zhang, D. Wang, Inorg. Chim. Acta 358 (2005) 4575–4580.[6] B. Setlow, P. Setlow, Appl. Environ. Microbiol. 59 (1993) 640–651.[7] J.T. Groves, I.O. Kady, Inorg. Chem. 32 (1993) 3868–3872.[8] D.C. Crans, M. Yang, T. Jakusch, T. Kiss, Inorg. Chem. 39 (2000) 4409–4416.[9] L.J.K. Boerner, J.M. Zaleski, Curr. Opin. Chem. Biol. 9 (2005) 135–144;

C.X. Zhang, S.J. Lippard, Curr. Opin. Chem. Biol. 7 (2003) 481–489.10] M. Goldstein, J.K. Barton, J.M. Goldberg, C.V. Kumar, N.J. Turro, J. Am. Chem. Soc.

10 (1986) 2081–2088.11] S. Delaney, M. Pascaly, P.K. Bhattacharya, K. Han, J.K. Barton, Inorg. Chem. 41

(2002) 1966–1974;M. Goldstein, J.K. Barton, H.M. Berman, Inorg. Chem. 25 (1986) 842–847.

12] P.J. Dardlier, R.E. Holmlin, J.K. Barton, Science 275 (1997) 1465–1468.13] F.R. Keene, Coord. Chem. Rev. 166 (1997) 121–154.

14] S. Delamey, M. Pascaly, P.K. Bhattacharya, K. Han, J.K. Barton, Inorg. Chem. 41

(2002) 1966–1974.15] H.T. Chifotides, K.R. Dunbar, Acc. Chem. Res. 38 (2005) 146–156.16] C. Tu, Y. Shao, N. Gan, Q. Xu, Z.J. Guo, Inorg. Chem. 43 (2004) 4761–4766.17] H. Sakurai, Y. Koyima, Y. Yoshikawa, K. Kawabe, H. Yasui, Coord. Chem. Rev. 226

(2002) 187–211, and references therein.

[

[

[

20.5 16.4 15.9 15.3 14.823.4 18.2 17.0 18.3 18.125.5 19.8 18.6 20.3 20.130.0 20.0 22.0 22.0 22.0

18] A.C. Gozalez-Baro, E.E. Castellano, O.E. Piro, B.S. Parajon-Costa, Polyhedron 24(2005) 49–55.

19] M. Kathryn, T. Deck, A. Tseng, J.N. Burstyn, Inorg. Chem. 41 (2002) 669–677.20] L.-Z. Li, C. Zhao, T. Xu, H.-W. Ji, Y.-H. Yu, G.-Q. Guo, H. Chao, J. Inorg. Biochem.

99 (2005) 1076–1082.21] K. Dhara, J. Ratha, M. Manassero, X.-Y. Wang, S. Gao, P. Banerjee, J. Inorg. Chem.

101 (2007) 95–103.22] C.N. Reilly, R.W. Schmid, F.S. Sadek, J. Chem. Educ. 36 (1959) 619–626.23] K.R. Rupesh, S. Deepalatha, M. Krishnaveni, R. Venkatesan, S. Jayachandran, Eur.

J. Med. Chem. 41 (2006) 1494–1503.24] J. Sambrook, E.F. Fritsch, T. Maniatis, Molecular Cloning: A Laboratory Manual,

2nd ed., Cold Spring Harbor Laboratory Press, Cold Spring Harbor, 1989.25] W.J. Geary, Coord. Chem. Rev. 7 (1971) 81–122.26] G.B. Deacon, R.J. Philips, Coord. Chem. Rev. 33 (1980) 227–250.27] M. Devereux, M. McCann, V. Leon, V. McKee, R.J. Ball, Polyhedron 21 (2002)

1063–1071.28] K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Com-

pounds, 3rd ed., Wiley, New York, 1978.29] A. Ciccarese, D.A. Clemente, F.P. Fanizzi, A. Marzotto, G. Valle, Acta Cryst C54

(1998) 1779–1781.30] P.W. Wade, R.D. Hancock, C.A. Boeyens, S.M. Dobson, J. Chem. Soc. Dalton Trans.

(1990) 483–488.31] C. Ma, J. Lee, R. Zhang, D. Wang, Inorg. Chim. Acta 358 (2005) 4575–4580.32] B.A. Prakasam, K. Ramalingam, M. Saravanam, G. Bocelli, A. Cantoni, Polyhedron

23 (2004) 77–82.33] E.K. Efthimiadou, Y. Sanakis, N. Katsaros, A. Karaliota, G. Psomas, Polyhedron

26 (2007) 1148–1158.34] K.S. Siddiqi, Sadaf Khan, Shahab A.A. Nami, J. Inclusion Phenom. Macrocyclic

Chem. 55 (2006) 359–366.35] F.A. Cotton, Progress in Inorganic Chemistry, vol. 6, Interscience Publishers,

U.S.A., 1964.36] E.Q. Gao, S. Bi, H. Sun, S. Liu, Synth. React. Inorg. Met. Org. Chem. 27 (1997) 1115.37] B.N. Figgis, Introduction to Ligand Fields, Wiley Eastern Ltd., New Delhi, 1976.38] A.B.P. Lever, Electronic Spectra of Inorganic and Coordination Compounds, John

Wiley, New York, 1968.39] M. Vicente, C. Lodeiro, H. Adams, R. Bastida, A. Blas, D.E. Fenton, A. Macyas, A.

Rodryguez, T. Rodryguez-Blas, Eur. J. Inorg. Chem. (2000) 1015–1024.40] M.J. Prushan, A.W. Anthony, R.J. Butcher, Inorg. Chim. Acta 992 (2000) 300–302.41] F.A. Cotton, G. Wilkinson, Advanced Inorganic Chemistry, John Wiley & Sons,

New York, 1988.42] M.S.S. Babu, K.H. Reddy, P.G. Krishna, Polyhedron 26 (2007) 572–580.43] P.M. Bush, J.P. Whitehead, C.C. Pink, E.C. Gramm, J.L. Eglin, S.P. Watton, L.E. Pence,

Inorg. Chem. 40 (2001) 1871–1877, and references therein.44] I. Ucar, et al., J. Mol. Struct. 834–836 (2007) 336–344.45] M. Chatterjee, M. Maji, S. Ghosh, T.C.W. Mak, J. Chem. Soc. Dalton Trans. (2000)

2875–2883.46] T. Gupta, A.K. Patra, S. Dhar, M. Nethaji, A.R. Chakravarty, J. Chem. Sci. 117 (2005)

123–132.47] S.P. Sovilj, K. BabicÂ-SamardzÏija, D.M. Minic, Thermochim. Acta 370 (2001)

29–36.48] I.T. Ahmed, Spectrochim, Acta Part A 65 (2006) 5–10.49] E. Nyarko, N. Hanada, A. Habib, M. Tabata, Inorg. Chim. Acta 357 (2004) 739–745.

50] Y. An, S.-D. Liu, S.-Yi Deng, L.-N. Ji, Z.-Wan Mao, J. Inorg. Biochem. 100 (2006)

1586–1593.51] G. Psomas, A. Tarushi, E.K. Efthimiadou, Y. Sanakis, C.P. Raptopoulou, N. Kat-

saros, J. Inorg. Biochem. 100 (2006) 1764–1773.52] R.P. John, A. Sreekanth, V. Rajakamman, T.A. Ajith, M.R.P. Kurup, Polyhedron 23

(2004) 2549–2559.