Journal of Photochemistry & Photobiology, B: Biology 157 (2016) 77–88

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Journal of Photochemistry & Photobiology, B: Biology journal homepage: www.elsevier.com/locate/jphotobiol

Exploring the photochemosensitivity by novel cysteine-based mixed ligand complexes Muthusamy Selvaganapathy 1, Narayanaperumal Pravin, Vellaichamy Muniyandi, Mohammed Nazeer, Natarajan Raman ⁎ Research Department of Chemistry, VHNSN College, Virudhunagar 626 001, Tamil Nadu, India

a r t i c l e

i n f o

Article history: Received 7 April 2015 Accepted 3 February 2016 Available online xxxx Keywords: Cysteine Intercalative mode Hydrolytic mechanism Photocytotoxicity

a b s t r a c t A new series of cysteine-based metal(II) complexes with 2,2′-bipyridine or 1,10-phenanthroline as co-ligand have been prepared and characterized. Their DNA binding and cleavage properties have been studied. The analytical and spectroscopic data of complexes 1–18 reveal that the complexes adopt an octahedral geometry around the central metal ion in which the cysteine is coordinated through NS and NN atoms, respectively. Spectroscopic titration and viscosity measurements reveal that the complexes bind to DNA through an intercalative mode. Electrophoresis measurements exhibit that they cleave pBR322 DNA efficiently in the presence of 3mercaptopropionic acid (MPA), probably via hydrolytic mechanism with the involvement of •OH. The in vitro anticancer activities indicate that the Cu(II) complexes are active against four selected human tumor cell lines. Furthermore, it is remarkable that all the complexes exhibit significant photocytotoxicity against human breast cancer cell lines (MCF-7) with a potency more than the widely used drugs photofrin and cisplatin indicating that they have the potential to act as effective anticancer drugs in a dose-dependent manner. © 2016 Elsevier B.V. All rights reserved.

1. Introduction Cancer is a deadly global menace (with total number of global cancer cases in 2008 being 12.6 million; in 2030, this number is expected to increase by +69%). Cancer or malignant neoplasm is a broad group of various diseases, all involving unregulated cell growth, invasion, and metastasis [1]. The current treatment regime for cancer primarily includes surgery and chemotherapy; however, chemotherapy is used as a mainstay due to its ability to cure widespread malignancies or metastatic cancers either alone or in combination with radiotherapy [2]. Nevertheless, the curative effects of the existing chemotherapeutic drugs are not good enough owing to severe health side effects, acquisition of resistance by tumor cells, and cost factors. Within the world of clinical oncology, the chemoprevention of cancer is perceived to be a failure. One major obstacle is that cancers evolve from numerous tissues (different phenotypes) with multiple etiologies and endless combinations of genetic or epigenetic alterations, and therefore, the one-size-fits-all therapy approach cannot be undertaken. Photodynamic therapy (PDT) drugs are activated by light to achieve spatiotemporally controlled chemotherapeutic action [3]. It provides a means to circumvent the drawbacks of conventional cancer therapies, most importantly through low systemic toxicity, localized action to ⁎ Corresponding author. E-mail address: [email protected] (N. Raman). 1 Present address: Department of Chemistry, Kodaikanal Institute of Technology, Machur, Kodaikanal-624,104, Tamil Nadu, India.

http://dx.doi.org/10.1016/j.jphotobiol.2016.02.008 1011-1344/© 2016 Elsevier B.V. All rights reserved.

irradiated areas, and low level of invasiveness [4]. It is now recognized as an alternative and in some cases a superior approach to conventional treatments in dermatology and for endoscopically accessible tumors. These include bladder, gastrointestinal, esophageal, prostate, and gynecological lesions [4,5], in both early- and late-stage head and neck cancers [4,6] and in inoperable early central lung cancers [7]. Therefore, it is imperative for chemists to develop ideal anticancer agents not only with good water solubility and accessible clinical value but also bringing fewer side effects, preferably non-covalently binding. Inorganic complexes have also been investigated for use in PDT, and transition metal complexes which undergo photoinduced ligand exchange represent one class of potential photochemotherapeutic (PCT) agents that do not require O2 for activity [8,9]. Following light absorption, these agents can deliver caged ligands that become cytotoxic upon their release [10,11], or the resulting metal photoproduct may preferentially bind biomolecules following photoinduced aquation [12]. Moreover, recently, many research groups have reported the in vitro biological potential of novel transition metal complexes and to elucidating the anticancer potential of these complexes with bipyridine (bpy) or phenantholine (phen) heterocyclic ligands [13–16]. Cysteine is known to act as an active site in bio-performance of the enzymes known as cysteine protease. Transition metal–cysteine complexes have considerable biological activities [17] and some of their antitumor properties show promise in therapeutic applications [18]. It is eminent that photoactivated chemotherapy (PACT) provides control over when and where a drug is activated, resulting in a greater specificity of drug action, less side effects, and thus it has remarkable potential

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M. Selvaganapathy et al. / Journal of Photochemistry & Photobiology, B: Biology 157 (2016) 77–88

for the treatment of cancer [19]. An inactive pro-drug is activated in the cell upon irradiation and the specificity of the drug is increased by the minimization of its toxicity in the surrounding healthy tissue [20]. The development of new metal complexes with high cytotoxicity which could be enhanced upon irradiation provides a highly challenging goal in the field of bioinorganic chemistry. Therefore, the reactivity of metal complexes toward DNA is useful in the design and synthesis of metal-based anticancer therapeutics. In this respect, in our studies on the development of new metallotherapeutics, we have synthesized a series of novel cysteine Schiff bases and their corresponding metal(II) complexes using phen or bpy as co-ligand and assessed their interaction with DNA. The cytotoxicity of the complexes has been assessed using four human cancer cell lines under in vitro condition. This serves to highlight the clinical potential of this series of compounds. We hope that the obtained results may contribute to the rational molecular design of DNA targeting reagents with high affinity and specificity as potential antitumor chemotherapeutic agents, as well as elucidate valuable information to understand their specific delivery at the active site of action, besides providing the pharmacological behaviors in vitro.

2. Experimental Protocol The materials and methods for the pharmacological experimental procedures were reported in our previous paper [21]. The powder X-ray diffraction, ESI-Mass and fluorescence spectral methodologies are given in S1 (Supplementary files). The Schiff base ligands, L1–L3 (condensation of cysteine and p-substituted benzaldehyde, L1 = −NO2, L2 = −H, and L2 = −OCH3) and their mixed-ligand (bpy or phen) Cu(II),Co(II), and Zn(II) complexes were prepared by the following procedure.

2.1. Synthesis of Schiff Base Ligands Cysteine (2.04 g, 0.01 mol) was dissolved in 20 mL of a water– ethanol mixture (1:1) and the solution was stirred to obtain a homogeneous solution. Then, an ethanolic solution of substituted benzaldehyde (0.01 mol) was added to this solution dropwise, and the resultant mixture was refluxed for ca. 5 h. A pale yellow-colored solution was obtained, which was reduced to one-third on a water bath. The resultant crystalline product that precipitated was filtered off, washed, and recrystallized with cold methanol and finally dried in vacuum. Yield 77–86%. L1: Yield: 77%. F.W: 254; m.p: N 180 °C; FT-IR (KBr): 1638 (HC = N), 1466 tasy(COO), 1375 tsy(COO), 2739 (\\SH) cm−1; 1H NMR (DMSO-d6): 8.9 (s, 1 H, CH = N), 7.3–8.0 (m, 5 H, aromatic C), 4.4 (t, 1 H, CH) 11.0 (s, 1 H, COOH), 3.8 (s, 3 H, O\\CH3), 3.0–3.3 (t, 2 H, S\\CH2) ppm; 13 C NMR (DMSO-d6 ): δ = 71.9 (CH), 177.5 (COOH), 122.6–147.2 (aromatic C), 162.3 (CH_N), 26.4 (S\\CH2) ppm; MS: m/z = 254, λmax in DMSO 42,564, 36,978 cm− 1. L2: Yield: 81%. F.W: 209; m.p: N 180 °C; FT-IR (KBr): 1624 (HC_N), 1460 t asy(COO), 1371 tsy (COO), 2748 (\\SH) cm− 1 ; 1 H NMR (DMSO-d 6 ): 9.1 (s, 1 H, CH_N), 7.4–7.9 (m, 5 H, aromatic C), 4.4 (t, 1 H, CH) 11.0 (s, 1 H, COOH), 3.8 (s, 3 H, O\\CH3), 3.0–3.2 (t, 2 H, S\\CH2) ppm; 13C NMR (DMSO-d6): δ = 71.8 (CH), 177.5 (COOH), 122.8–147.0 (aromatic C), 161.7 (CH_N), 26.4 (S-CH2) ppm; MS: m/z = 209, λmax in DMSO 42,424, 36,818 cm− 1. L3: Yield: 86%. F.W: 239; m.p: N 185 °C; FT-IR (KBr): 1616 (HC_N), 1457 tasy(COO), 1368 tsy(COO), 2762 (\\SH) cm−1; 1H NMR (DMSOd6): 9.0 (s, 1 H, CH_N), 7.4–8.0 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 11.0 (s, 1 H, COOH), 3.8 (s, 3 H, O\\CH3), 3.1–3.3 (t, 2 H, S\\CH2) ppm; 13C NMR (DMSO-d6): δ = 71.6 (CH), 177.3 (COOH), 122.6–147.2 (aromatic C), 162.4 (CH_N), 55.8 (O\\CH3), 26.3 (S\\CH2) ppm; MS: m/z = 239, λmax in DMSO 42,536, 36,783 cm−1.

2.2. Synthesis of Metal Complexes The complexes were prepared by mixing the appropriate molar quantity of the ligand L1–L3 with the metal salts [Cu(II), Co(II), and Zn(II)] using the following procedure: an ethanolic solution of ligand (0.003 mol) was stirred with 5 mL of an ethanolic solution of the anhydrous metal(II) chloride (0.003 mol) for ca. 1 h. A methanolic solution (5 mL) of 2,2′-bipyridine or 1,10-phenanthroline (0.006 mol) was added to this mixture, and the stirring was continued for 1 h. The solid product obtained was filtered and washed with ethanol. [CuL1(bpy)2]Cl. (1) Yield: 78%. F.W: 664; m.p: N219 °C; Anal. Calc. for C30H25ClCuN6O4S: C, 54.2; H, 3.8; Cu, 9.6; N, 12.6%. Found: C, 54.1; H, 3.6; Cu, 9.4; N, 12.3%. FT-IR (KBr): 1615 (HC_N), 1457 υasy(COO−), 1369 υsy(COO−); 465 (M − N) 378 (M − S) cm−1; Λm × 10− 3 (Ω−1 mol−1 cm2) 49.7; μeff (BM) 1.84; MS: m/z = 630 (M + H), λmax in DMSO 31,485, 27,143, 13,671 cm−1 (ε = 65 L M−1 cm−1). [CuL2(bpy)2]Cl. (2) Yield: 75%. F.W: 619; m.p: N 215 °C; Anal. Calc. for C30H26ClCuN5O2S: C, 58.2; H, 4.2; Cu, 10.3; N, 11.3%. Found: C, 58.0; H, 4.1; Cu, 10.1; N, 11.1%. FT-IR (KBr): 1619 (HC_N), 1466 υasy(COO−), 1375 υsy(COO−); 461 (M − N) 382 (M − S) cm−1; Λm × 10− 3(Ω−1 mol−1 cm2) 50.3; μeff (BM) 1.83; MS: m/z = 585 (M + H), λmax in DMSO 32,825, 27,186, 13,652 cm− 1 (ε = 68 L M−1 cm−1). [CuL3(bpy)2]Cl. (3) Yield: 73%. F.W: 649; m.p: N 223 °C; Anal. Calc. for C31H28ClCuN5O3S: C, 57.3; H, 4.3; Cu, 9.8; N, 10.8%. Found: C, 57.1; H, 4.1; Cu, 9.6; N, 10.6%. FT-IR (KBr): 1628 (HC = N), 1459 υasy(COO−), 1371 υsy(COO−); 455 (M − N) 375 (M − S) cm−1; Λm × 10−3 (Ω−1 mol−1 cm2) 48.5; μeff (BM) 1.82; MS: m/z = 615 (M + H), λmax in DMSO 32,457, 27,236, 13,656 cm−1 (ε = 63 L M−1 cm−1). [CoL1(bpy)2]Cl. (4) Yield: 75%. F.W: 660; m.p: N219 °C; Anal. Calc. for C30H25ClCoN6O4S: C, 54.6; H, 3.8; Co, 8.9; N, 12.7%. Found: C, 54.5; H, 3.6; Co, 8.7; N, 12.5%. FT-IR (KBr): 1622 (HC_N), 1462 υasy(COO−), 1371 υsy(COO−); 471 (M − N) 385 (M − S) cm−1; Λm × 10− 3 (Ω−1 mol−1 cm2) 53.4; μeff (BM) 4.86; MS: m/z = 625 (M + H), λmax in DMSO 23,045, 17,491, 14,891 cm−1 (ε = 48 L M−1 cm−1). [CoL2(bpy)2]Cl. (5) Yield: 72%. F.W: 615; m.p: N210 °C; Anal. Calc. for C30H26ClCoN5O2S: C, 58.6; H, 4.3; Co, 9.6; N, 11.4%. Found: C, 58.5; H, 4.1; Co, 9.3; N, 11.2%. FT-IR (KBr): 1613 (HC_N), 1459 υasy(COO−), 1369 υsy(COO−); 480 (M − N) 381 (M − S) cm−1; Λm × 10− 3 (Ω−1 mol−1 cm2) 54.9; μeff (BM) 4.82; MS: m/z = 580 (M + H), λmax in DMSO 22,715, 17,523, 14,679 cm−1 (ε = 54 L M−1 cm−1). [CoL3(bpy)2]Cl. (6) Yield: 76%. F.W: 645; m.p: N219 °C; Anal. Calc. for C31H28ClCoN5O3S: C, 57.7; H, 4.4; Co, 9.1; N, 10.9%. Found: C, 57.5; H, 4.2; Co, 9.0; N, 10.7%. FT-IR (KBr): 1598 (HC_N), 1455 υasy(COO−), 1370 υsy(COO−); 473 (M − N) 385 (M − S) cm−1; Λm × 10− 3 (Ω−1 mol−1 cm2) 53.4; μeff (BM) 4.85; MS: m/z = 610 (M + H), λmax in DMSO 22,753, 17,610, 14,642 cm−1 (ε = 52 L M−1 cm−1). [ZnL1(bpy)2]Cl. (7) Yield: 68%. F.W: 666; m.p: N215 °C; Anal. Calc. for C30H25ClN6O4SZn: C, 54.1; H, 3.8; N, 12.6; Zn, 9.8%. Found: C, 54.0; H, 3.6; N, 12.5; Zn, 9.6%. FT-IR (KBr): 1622 (HC_N), 1462 υasy(COO−), 1371 υsy(COO−); 472 (M − N) 386 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.6 (s, 1 H, CH_N), 7.6–8.7 (m, 16 H, bpy), 7.5–8.0 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 10.8 (s, 1 H, COOH), 3.0–3.2 (t, 2 H, S\\CH2); 13C NMR (ppm) (DMSO-d6): 71.7 (CH), 177.3 (COOH), 124.6–148.2 (aromatic C), 120.6–146.1 (bpy), 158.3 (CH_N), 26.4 (S\\CH2); Λm × 10−3(Ω−1 mol−1 cm2) 55.4, MS: m/z = 633 (M+). [ZnL2(bpy)2]Cl. (8) Yield: 68%. F.W: 621; m.p: N210 °C; Anal. Calc. for C30H26ClN5O2SZn: C, 57.9; H, 4.2; N, 11.3; Zn, 10.5%. Found: C, 57.7; H, 4.0; N, 11.2; Zn, 10.3%. FT-IR (KBr): 1627 (HC_N), 1462 υasy(COO−), 1371 υsy(COO−); 474 (M − N) 389 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.7 (s, 1 H, CH_N), 7.5–8.7 (m, 16 H, bpy), 7.4–7.8 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 10.8 (s, 1 H, COOH), 3.0–3.2 (t, 2 H, S\\CH2); 13C NMR (ppm) (DMSO-d6): 71.8 (CH), 177.4 (COOH), 124.6–148.3 (aromatic C), 120.6–146.2 (bpy), 159.2 (CH_N), 26.4 (S\\CH2); Λm × 10− 3(Ω− 1 mol−1 cm2) 59.7, MS: m/z = 588 (M+).

M. Selvaganapathy et al. / Journal of Photochemistry & Photobiology, B: Biology 157 (2016) 77–88

[ZnL3(bpy)2]Cl. (9) Yield: 68%. F.W: 651; m.p: N215 °C; Anal. Calc. for C31H28ClN5O3SZn: C, 57.1; H, 4.3; N, 10.7; Zn, 10.0%. Found: C, 57.0; H, 4.1; N, 10.6; Zn, 9.9%. FT-IR (KBr): 1622 (HC_N), 1466 υasy(COO−), 1375 υsy(COO−); 476 (M − N) 387 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.6 (s, 1 H, CH_N), 7.6–8.5 (m, 16 H, bpy), 7.0–7.7 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH), 10.8 (s, 1 H, COOH), 3.8 (s, 3 H, O-CH 3 ), 3.0–3.2 (t, 2 H, S\\CH 2 ); 13 C NMR (ppm) (DMSO-d 6 ): 71.9 (CH), 177.5 (COOH), 124.5–148.2 (aromatic C), 120.4–146.1 (bpy), 158.8 (CH_N), 55.8 (O\\CH 3 ), 26.4 (S\\CH 2 ) ppm; Λm × 10− 3(Ω− 1 mol− 1 cm2) 60.5, MS: m/z = 618 (M +). [CuL1(phen)2]Cl. (10) Yield: 73%. F.W: 712; m.p: N240 °C; Anal. Calc. for C34H26ClCuN5O2S: C, 57.3; H, 3.5; Cu, 8.9; N, 11.8%. Found: C, 57.1; H, 3.3; Cu, 8.7; N, 11.5%. FT-IR (KBr): 1606 (HC_N), 1460 υasy(COO−), 1372 υsy (COO−); 459 (M − N) 384 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 56.9; μeff (BM) 1.87; MS: m/z = 678 (M + H), λmax in DMSO 32,856, 27,293, 13,752 cm− 1 (ε = 69 L M− 1 cm− 1). [CuL2(phen)2]Cl. (11) Yield: 71%. F.W: 667; m.p: N235 °C; Anal. Calc. for C34H26ClCuN5O2S: C, 61.2; H, 3.9; Cu, 9.5; N, 10.5%. Found: C, 61.0; H, 3.7; Cu, 9.3; N, 10.4%. FT-IR (KBr): 1609 (HC_N), 1466 υasy(COO−), 1378 υsy (COO−); 472 (M − N) 386 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 59.3; μeff (BM) 1.86; MS: m/z = 633 (M + H), λmax in DMSO 33,205, 27,381, 13,671 cm− 1 (ε = 67 L M− 1 cm− 1). [CuL3(phen)2]Cl. (12) Yield: 76%. F.W: 697; m.p: N250 °C; Anal. Calc. for C35H28ClCuN5O3S: C, 60.2; H, 4.0; Cu, 9.1; N, 10.0%. Found: C, 60.0; H, 3.9; Cu, 9.0; N, 9.9%. FT-IR (KBr): 1596 (HC_N), 1461 υasy(COO−), 1374 υsy (COO−); 484 (M − N) 390 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 62.9; μeff (BM) 1.89; MS: m/z = 663 (M + H), λmax in DMSO 33,157, 27,576, 13,826 cm− 1 (ε = 72 L M− 1 cm− 1). [CoL1(phen)2]Cl. (13) Yield: 75%. F.W: 708; m.p: N230 °C; Anal. Calc. for C34H25ClCoN6O4S: C, 57.7; H, 3.6; Co, 8.3; N, 11.8%. Found: C, 57.5; H, 3.5; Co, 8.1; N, 11.6%. FT-IR (KBr): 1612 (HC_N), 1469 υasy(COO−), 1373 υsy (COO−); 480 (M − N) 388 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 57.6; μeff (BM) 4.89; MS: m/z = 673 (M + H), λmax in DMSO 22,946, 17,573, 14,642 cm− 1 (ε = 51 L M− 1 cm− 1). [CoL2(phen)2]Cl. (14) Yield: 74%. F.W: 663; m.p: N210 °C; Anal. Calc. for C34H26ClCoN5O2S: C, 61.6; H, 3.9; Co, 8.9; N, 10.5%. Found: C, 61.4; H, 3.7; Co, 8.7; N, 10.4%. FT-IR (KBr): 1618 (HC_N), 1462 υasy(COO−), 1369 υsy (COO−); 484 (M − N) 386 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 54.3; μeff (BM) 4.84; MS: m/z = 628 (M + H), λmax in DMSO 22,836, 17,593, 14,759 cm− 1 (ε = 50 L M− 1 cm− 1). [CoL3(phen)2]Cl. (15) Yield: 78%. F.W: 693; m.p: N255 °C; Anal. Calc. for C35H28ClCoN5O3S: C, 60.6; H, 4.0; Co, 8.5; N, 10.1%. Found: C, 60.3; H, 3.8; Co, 8.3; N, 10.0%. FT-IR (KBr): 1597 (HC_N), 1460 υasy(COO−), 1371 υsy (COO−); 479 (M − N) 387 (M − S) cm − 1 ; Λm × 10− 3(Ω− 1 mol− 1 cm2) 58.1; μeff (BM) 4.92; MS: m/z = 658 (M + H), λmax in DMSO 22,839, 17,698, 14,534 cm− 1 (ε = 59 L M− 1 cm− 1). [ZnL1(phen)2]Cl. (16) Yield: 74%. F.W: 663; m.p: N210 °C; Anal. Calc. for C34H26ClCoN5O2S: C, 61.6; H, 3.9; Co, 8.9; N, 10.5%. Found: C, 61.4; H, 3.7; Co, 8.7; N, 10.4%. FT-IR (KBr): 1618 (HC_N), 1462 υasy(COO−), 1369 υsy(COO−); 484 (M − N) 386 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.7 (s, 1 H, CH_N), 7.6–8.6 (m, 16 H, phen), 7.3–7.7 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 10.8 (s, 1 H, COOH), 3.0–3.2 (t, 2 H, S\\CH2); 13C NMR (ppm) (DMSO-d6): 71.8 (CH), 177.2 (COOH), 124.8–148.6 (aromatic C), 121.5–146.7 (phen), 159.2 (CH_N), 26.4 (S\\CH2); Λm × 10−3(Ω−1 mol−1 cm2) 54.3, MS: m/z = 681 (M+). [ZnL2(phen)2]Cl. (17) Yield: 70%. F.W: 669; m.p: N200 °C; Anal. Calc. for C34H26ClN5O2SZn: C, 61.0; H, 3.9; N, 10.4; Zn, 9.7%. Found: C, 59.7; H, 3.6; N, 10.2; Zn, 9.4%. FT-IR (KBr): 1610 (HC_N), 1462 υasy(COO−), 1369 υsy(COO−); 480 (M − N) 382 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.8 (s, 1 H, CH_N), 7.6–8.4 (m, 16 H, phen), 7.3–7.7

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(m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 10.8 (s, 1 H, COOH), 3.0–3.2 (t, 2 H, S\\CH2); 13C NMR (ppm) (DMSO-d6): 71.8 (CH), 177.5 (COOH), 124.8–148.6 (aromatic C), 121.5–146.7 (phen), 159.2 (CH_N), 26.4 (S\\CH2); Λm × 10−3(Ω−1 mol−1 cm2) 61.3, MS: m/z = 636 (M+). [ZnL3(phen)2]Cl. (18) Yield: 67%. F.W: 699; m.p: N 230 °C; Anal. Calc. for C35H28ClN5O3SZn: C, 60.1; H, 4.0; N, 10.0; Zn, 9.3%. Found: C, 60.0; H, 3.9; N, 9.8; Zn, 9.1%. FT-IR (KBr): 1618 (HC_N), 1465 υasy(COO−), 1370 υsy(COO−); 478 (M − N) 380 (M − S) cm− 1; 1H NMR (ppm) (DMSO-d6): 8.9 (s, 1 H, CH_N), 7.6–8.5 (m, 16 H, phen), 7.2–7.6 (m, 5 H, aromatic C), 4.3 (t, 1 H, CH) 10.7 (s, 1 H, COOH), 3.8 (s, 3 H, O\\CH3), 3.0–3.2 (t, 2 H, S\\CH 2 ); 13 C NMR (ppm) (DMSO-d 6 ): 71.8 (CH), 177.3 (COOH), 124.5–148.3 (aromatic C), 121.5–146.6 (phen), 158.4 (CH_N), 55.8 (O\\CH 3 ), 26.4 (S\\CH 2 ) ppm; Λm × 10− 3(Ω− 1 mol− 1 cm2) 62.9, MS: m/z = 666 (M +). 3. Results and Discussion Eighteen mixed-ligand complexes [ML(bpy/phen)2]Cl (where M is Cu(II), Co(II), and Zn(II), L is p-substituted benzaldehyde analogues (L1 = NO2/L2 = − H/L3 = − OCH3) and bpy/phen is 2,2′-bipyridine/ 1,10-phenanthroline) have been prepared by the reaction of metal salts and Schiff base with bpy/phen ligands (Scheme 1). They have been characterized by elemental analysis, IR, UV–vis, 1H-NMR, 13 C-NMR, ESI-MS, and EPR spectroscopic methods. 3.1. Elemental Analysis vs Molar Conductance The elemental analysis data of the Schiff base ligands (L1–L3) and their metal complexes agree well with the assigned formulae of the proposed structure (Scheme. 1). The synthesis, spectroscopic and analytical data of prepared complexes are presented in the Experimental section. The metal complexes were dissolved in DMSO and the molar conductivities of 10−3 mol/dm3 of their solution at 25 °C were measured. The higher molar conductance values (48.5–62.9 Ω−1 cm−2 mol−1) of the complexes support their electrolytic nature. The relative high molar conductivities show that these complexes are stable in solution. Both the ligands and their complexes are very stable at room temperature in the solid as well as solution state. 3.2. Magnetic vs Electronic Behavior The UV–vis spectra of the complexes were recorded in DMSO solution. All the complexes show the high energy absorption band in the region 27,143–33,216 cm− 1. This transition may be attributed to the charge transfer band. The electronic spectra of Cu(II) complexes display the d–d transition bands in the region 13,648–13,847 (ε = 63– 72 L M−1 cm−1) cm−1 which are due to 2Eg → 2T2g transition. These d–d transition bands strongly favor a distorted octahedral geometry around the metal ion. Their magnetic susceptibility values (1.82–1.89 B.M.) are typical for mononuclear Cu(II) compound having d9electronic configuration. The electronic spectra of Co(II) complexes display three d–d transition bands in the region 14,879–14,534 (ε = 48–59 L M−1 cm−1), 17,491–17,698, and 22,628–23,045 cm−1 which are assigned to 4T1g(F) → 4T2g(F) (υ1), 4T1g(F) → 4A2g(F) (υ2), and 4 T1g(F) → 4T2g(P) (υ3) transitions, respectively. This indicates that the Co(II) complexes are six coordinate and probably an octahedral geometry, which is also supported by their magnetic susceptibility values in the range 4.81–4.92 B.M. 3.3. IR Spectra The IR spectra of the complexes were very similar to each other, except some slight shifts and intensity change of a few vibration bands caused by different metal ions, which indicate that the complexes have similar structures. In the infrared spectra, the band of the Schiff base ligands (L1–L3) observed at 1638–1616 cm− 1 was shifted to

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Scheme 1. The schematic representation of amino acid mixed-ligand complexes.

lower frequency by 10–33 cm−1 on complexation, suggesting the coordination of the azomethine nitrogen. A characteristic strong band at 2739–2762 cm−1 is ascribed to υ(S\\H), disappeared in the spectra of metal complexes, confirming deprotonation and coordination of thiol group [22]. This is further supported by the lower frequency band appeared at 750–765 cm− 1 in the metal complexes due to υ(C\\S). The broad bands corresponding to free ligands at 1375–1368 and 1466–1457 cm− 1 attributed to the carbonyl bonds υsy(COO−) and υasy(COO−), did not show any significant shift on complex formation, thus denoting the non-participation of carbonyl oxygen in the complexation. This evidence indicates that sulphur plays a major role in binding and carboxylate plays a minor role. Moreover, the appearance of new peaks is also guiding for chelation. The new bands in the region of 375–390 and 455–484 cm− 1 in all metal complexes are assigned to vibrations of υ(M\\S) and υ(M\\N) bonds, respectively [23]. 3.4. Nuclear Magnetic Resonance Spectra The 1H-NMR spectra of the Schiff bases (L1–L3) and their zinc(II) complexes were recorded at room temperature in DMSO-d6. The 1H NMR spectra of the ligand L1 and its Zn(II) complex 16 are shown in Fig.S1. The ligands (L1–L3) showed one singlet at 8.9–9.1 δ which is due to azomethine (CH_N) proton and a phenyl multiplet signal at

7.3–8.0 ppm, due to aromatic protons of substituted benzaldehyde moiety present in the cysteine derived Schiff bases. In the 1H NMR spectra of complexes 7 and 16, the protons of Schiff base ligands were shifted to down field due to the coordination with metal ion. Upon comparison with the free ligands, the signal observed at 14 ppm can be assigned to the\\SH protons [23]. This signal disappears in the spectra of Zn(II) complexes, which confirms the coordination of ligand to metal ion through the deprotonated thiol group. Further, a set of multiplets were observed in the range δ 7.2–8.0 ppm due to the presence of aromatic protons in complexes of 7 and 16. The resonance peaks observed in the spectra of the complexes 7 and 16 at around 7.5–8.7 ppm were assigned to the bpy and phen protons. The azomethine proton (\\CH_N) signal in the spectra of the zinc complexes was shifted to down field compared to the free ligand, suggesting deshielding of azomethine group due to the coordination with metal ion. The 13C NMR spectra of ligands L1–L3 showed aromatic carbons at 124.5–148.6 ppm and they also showed \\COOH carbons at 177.3– 177.8 for ligands L1–L3. The signal observed at 26.4 ppm for S\\CH2 carbon in ligands and their complexes. Moreover the (HC_N) carbon at 162.6–162.1 ppm for ligands L1–L3 was shifted to down fields (158.3– 159.2) upon coordination indicating that (HC_N) group is participating in the complex formation. There are no appreciable changes in all other peaks.

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3.5. Mass Spectra Mass spectra of all the compounds stand in good agreement with proposed structure. The ESI-Mass spectra of the ligand (L1) and its Cu(II) complex 10 are presented in Fig. S2. The mass spectrum of Schiff base ligand L1 showed peak at m/z 254 corresponding to [C10H9N2O4S] ion. Moreover, the spectrum of L1 exhibited peaks for the fragments at m/z 207, 149, 122, and 76 corresponding to [C9H7N2O4] (M +), [C7H5N2O2] (M+), [C6H4NO2] (M+), and [C6H4] (M+), respectively. The spectra of complexes 1–3 showed molecular ion peaks at m/z 630 (M + H), 585 (M + H), and 615 (M + H), respectively. The spectra of complexes 10–12 showed molecular ion peaks at m/z 678 (M + H), 633 (M + H), and 663 (M + H), respectively. The complex 1 gave a fragment ion peak with loss of one fragment of phenanthroline peak at m/z 498 (M+). 3.6. Powder X-Ray Diffraction Studies Single crystal X-ray crystallographic investigation is the most precise source of information regarding the structure of the complexes, but the difficulty of obtaining crystalline complexes in proper symmetric form has rendered the powder X-ray diffraction method for such study. During powder X-ray diffraction analysis, complexes 1 and 10 exhibited sharp peaks with crystalline nature peaks which are shown in Fig. S3, while no peaks obtained for rest of the complexes demonstrating their amorphous nature [24]. Powder X-ray diffraction analysis is useful to conclude the structure, particle size of the synthesized mixed ligand complexes. From the diffractograms, it is evident that the strong peaks appeared at 2θ = 17.25°, 21.84°, 24.13°, and 27.28° for complex 1 and 16.21°, 23.82°, 26.51°, 27.94°, 31.14°, 38.06°, and 43.54° for complex 10, respectively, in that order confirm the complex formation. The XRD pattern shows the crystalline nature of the complexes. The observed interplanar spacing values (“d” in A°) and the Miller indices (h k l) values for complexes 1 and 10 are 4.96–1.97 and [2 2 0]–[6 4 0], respectively. Experiential average grain sizes of the complexes 1 and 10 are established to be 41.58 and 52.46 nm, respectively, signifying that these mixed ligand complexes are in microcrystalline state which might confirm the proposed composition. All of these diffraction peaks in the XRD pattern indicate cubic crystalline structure of Cu(II) complexes. But, we are abortive to develop a single crystal of any of these mixed ligand complexes. Finally, the above data indicate that the coordination occurs through the nitrogen of the pby/phen ring, azomethine nitrogen atom, and sulphur present in amino acid moiety to give the structures as shown in Scheme 1. 3.7. Electron Paramagnetic Resonance Spectra The EPR spectra of all of the Cu(II) complexes were recorded in DMSO at 300 and 77 K. X-Band EPR spectra of Cu(II) complex 10 at room temperature (273 K) and at liquid nitrogen temperature (77 K) are given in Fig.S4. The spin Hamiltonian parameters calculated for the complexes are given in Table S1. The trend in the observed “g” values of Cu(II) complexes at room temperature, was g║ N g┴ N ge (2.0023). This trend provides an evidence of localization of the unpaired electron in dx22−y orbital. The Cu(II) complexes exhibited A║ = 152–161 N A┴ = 34–45; g║ = 2.24–2.28 N g┴ = 2.05–2.06 suggesting that the complexes are present in axially elongated octahedral geometry and the unpaired electron lies predominantly in the dx22−y orbital. The axial symmetry parameter G is defined as.

G¼

g║ −2:0023= g┴ −2:0023

According to Hathaway [25], if the G value is N 4, the exchange interaction is negligible, while a value is b 4 gives an indication for

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considerable exchange interaction in the complex. The axial symmetry parameter (G) of the reported Cu(II) complexes was N 4 (G = 4.29– 5.40) suggesting that the local tetragonal axes are aligned parallel or slightly misaligned and the unpaired electron is present in the dx2−y2 orbital. This result also indicates that the exchange coupling effects are not operative in the present Cu(II) complexes. The shapes of the spectra are consistent with octahedral geometry around the Cu(II) center in the complexes [26]. Molecular orbital coefficients, α2 (a measure of the covalency of the in-plane σ-bonding between copper 3d orbital and the ligand orbitals), β2 (in-plane π-bonding), and γ2 (out-of-plane π-bonding) were calculated by using the following equations [27]: 

α2 ¼ Ajj =0:036 þ gjj −2:0027 þ 3=7 g┴ −2:0027 þ 0:04   β2 ¼ gjj −2:0027 E=−−8λα2

γ2 ¼ g┴ −2:0027 E= −−2λα2 where λ = −828 cm−1 for the free copper ion and E is the electronic transition energy. If α2 = 1 indicates complete ionic character, whereas α2 = 0.5 denotes 100% covalent bonding, with the assumption of negligibly small values of the overlap integral. The observed values of α2 (0.67–0.74), β2 (0.92–0.98) and γ2 (0.68–0.79) indicate the significant covalent character to the M\\L bond. The lower value of α2 compared to β2 indicates that the in-plane σ-bonding is more covalent than the in-plane π-bonding. 3.8. Electronic Absorption Studies UV–visible absorption spectroscopy is a versatile and normally engaged method to determine the binding characteristics of metal complexes with DNA. In general, hyperchromism and hypochromism are the spectral features of DNA concerning changes of its double helix structure. Additionally, the existence of a red shift is indicative of the stabilization of the DNA duplex [28]. The experiment was carried out keeping the concentration of the Cu(II) complexes constant and varying the concentration of CT-DNA. The ligand-centered transitions at 428– 364 nm for complexes 1–18, respectively, were considered for the corresponding absorptivity changes upon the incremental addition of DNA. The interaction of complexes 1 and 10 with CT-DNA was followed by recording the UV–visible spectra of the system (Fig. 1). Fig. 1 shows the electronic absorption spectra of complexes 1 and 10 in the absence and presence of CT-DNA. For complex 1, upon the addition of DNA, the band around 428 nm shows hypochromism by about 29.7% and a significant bathochromic (red) shift of 7 nm. For complex 10, under the same conditions, upon the addition of DNA, the band at 382 nm exhibits hypochromism of 38.4% with a 3 nm minor bathochromic shift. The changes in the spectrum indicate that the complex possesses a strong affinity for CT-DNA. In order to compare quantitatively the binding strength of all the complexes, the intrinsic binding constant, Kb value of all the complexes (1–18) with DNA was calculated by monitoring the changes in the absorbance at 428–364 nm with an increasing concentration of DNA. The intrinsic binding constant (Kb) is a useful tool to monitor the magnitude of the binding strength of compounds with CT-DNA (Table S2). It can be determined by monitoring the changes in the absorbance in the IL band at the corresponding λmax with increasing concentration of DNA and is given by the ratio of the slope to the y intercept in plots of [DNA]/(εa − εf) versus [DNA]. The magnitudes of the intrinsic binding constants (Kb) were calculated to be 1.03 ± 0.5 × 105 M−1 − 6.25 ± 0.2 × 106 M−1 for all the test compounds (1–18), respectively revealing that the complexes 1–18 bind to DNA via an intercalative mode. These results are similar to those reported earlier for the intercalative mode of various metallointercalators [29]. From the results obtained, it has been found that all the complexes exhibit a good binding affinity

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energy “Δ G” of compound–DNA complex was calculated, using the following equation:   −1 ΔG ¼ –RT lnKb kJ mol

Free energies of all the compounds (1–18) were evaluated as negative values showing the spontaneity of compounds–DNA interaction (Table S2). However, results indicate that complex 10 binds to DNA more spontaneously compared to other compounds. 3.9. Fluorescence Spectroscopic Studies The hypochromism observed in the UV spectra of the CT-DNA-complex solutions of variable proportions allows DNA complex studies through fluorescent spectroscopy. The fluorescence spectral technique is an effective method to study metal interaction with DNA and the relative binding of these complexes to CT-DNA [30]. Ethidium bromide (EB) emits intense fluorescent light in the presence of DNA due to its strong intercalation between the adjacent DNA base pairs. Since EB intercalates DNA through interactions with the minor groove, the displacement of EB (quantified by fluorescence) by the titration of a compound is suggestive of an intercalative binding [30]. The study involves addition of the complexes to DNA pre-treated with EB and then measurement of the intensity of emission. The emission spectra at 586 nm and at 605 nm of the EB (2.0 μM) solutions which contain CT-DNA (20 μM) in the absence or presence of various concentrations of complexes 1 and 10 (0–210 μM) were recorded upon their excitation at 447 nm and at 486 nm (Fig. 2). The apparent binding constant (Kapp) was calculated using the equation [31]: K EB ½EB ¼ K app ½drug

Fig. 1. Absorption spectra of complexes 1 (a) and 10 (b) in Tris–HCl buffer (pH 7.2) at 37 °C in presence of increasing amount of DNA (0–150 μM).

to DNA. This is mainly due to the chelation of the metal with the ligand. It is to be noted that complex 10 exhibits a relatively higher binding constant compared to other complexes. On the other hand, the higher binding affinity of complex 10 to DNA than that of complex 1 may be due to the greater planar area of phen, thus resulting in higher hydrophobicity of phen than that of bpy. These data indicate that the extent of hypochromism commonly parallels the intercalative strength of the ligand, and moreover, a higher planar area, an extended π system, hydrophobicity, and aromaticity lead to deep penetration and hence more stacking within the base pairs of DNA. The obtained data in our experiment indicate that the size and the shape of the intercalated ligand have a significant effect on the strength of the DNA binding, and the most suitable intercalating ligand leads to the highest affinity of complexes with DNA. The different DNA-binding properties of the complexes 1–18 are due to the difference in the co-ligands as well as nature of the central metal ion, which leads to a greater binding affinity to DNA. Furthermore, the binding mode needs to be proved through some more experiments. Binding constants are a measure of compound–DNA complex stability while free energy indicates the spontaneity/non-spontaneity of compound–DNA binding. From the values of binding constant, Kb, free

where [drug] is the concentration of the complex at a 50% reduction of the fluorescence, KEt-Br = 107 M−1, and [EB] = 2.0 μM [31]. The concentration of the drug at a 50% reduction of the fluorescence is calculated from the diagram of I0/I vs the concentration of the complex [Q] (Fig. S5), where I0 and I are the fluorescence intensities of the CT-DNA in the absence and presence of complexes. The apparent binding constants Kapp calculated for Cu(II) complexes are 4.56 × 105 M− 1 − 3.15 × 106 M−1, suggesting an intercalative binding of both complexes (Table S2). Furthermore, the quenching data were analyzed according to the Stern–Volmer equation and the Kq value is obtained as a slope from the plot of I0/I vs [Q] [32] (Fig. S5). The quenching plots illustrate that the quenching of EB bound to CT-DNA by the complexes are in good agreement with the linear Stern–Volmer equation. KSV is the Stern–Volmer dynamic quenching constant and [Q] is the total concentration of the quencher (complexes 1–18). The linear Stern–Volmer equation is given as follows:

I 0 =I ¼ K q ½Q  þ 1;

The quenching constants (Kq = 1.68 × 105 M−1 − 3.54 × 105 M−1 (1–3 and 10–12), have been calculated from the plot of I0/I vs [Q]. These results suggest that complex 10 intercalates more strongly than complex 1. However, their pharmacodynamical, pharmacological, and toxicological properties should be further studied in vivo. 3.10. Electrochemical Activation Cyclic voltammetry (CV) is widely used for the evaluation of mode of action and binding strength of drug–DNA interaction. This technique

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Fig. 2. Emission spectra of EB bound to DNA in the presence of (a) 1 and (b) 10 ([EB] = 2.0 μM, [DNA] = 20 μM, [complex] = 0–210 μM). The arrow shows the intensity changing upon increasing complex concentration.

is predominantly useful for metal-based compounds due to their accessible redox states [33]. The major function of metal–biological compounds involves oxidation–reduction reactions in which metal containing biological molecules react directly with molecular oxygen to produce free radicals. Copper plays a pivotal role in cell physiology as a catalytic cofactor in the redox chemistry of mitochondrial respiration, iron absorption, free radical scavenging, and elastic cross-linking. The peak potential and peak current of the compound changes in the presence of DNA if the compound interacts with it. The variation in peak potential and peak current can be exploited for the determination of binding parameters, whereas the shift in peak potential can be used to ascertain the mode of interaction. Cyclic voltammetry of complexes 1–18 was performed with a glassy carbon electrode at a scan rate of 0.8 V s−1 in DMSO containing Tris–HCl buffer (0.1 M) as a supporting electrolyte. It can be seen from Fig. 3 (a and b) in the absence of DNA, the Cu(II) complexes 1 and 10 have been found to show a quasi-reversible redox process corresponding to Cu(II)/Cu(I) with the cathodic (Epc) and anodic peak potential (Epa) being − 0.268/0.066 and − 0.081/− 0.104 V (Δ Ep = − 0.187/ −0.107 V), respectively. The large peak-to-peak potential difference is suggestive of electrochemical reaction coupled with a chemical

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Fig. 3. Cyclic voltammograms of complexes 1 (a) and 10 (b) in Tris–HCl buffer (pH 7.2) at 37 °C in presence of increasing amount of DNA (0–120 μM).

reaction. With the increase in concentration of DNA in a constant amount of compound, the voltametric response of the compound changed as is evidenced by the sequential drop in peak current and gradual peak potential shift in positive direction. This behavior is suggestive of intercalative mode of interaction of investigated compounds with DNA. The observed decrease in peak current indicates the formation of large and slowly diffusing copper complex-DNA adduct due to which the free drug concentration (which is mainly responsible for the conduction of the current) is lowered indicating that the electron transfer process in the metal complex slows down in presence of DNA [34]. These results strongly demonstrate the interaction of Cu(II) complex with DNA. The voltammetic parameters obtained for the compounds without and in the presence of DNA are given in Table S3. Similarly, Table S3 describes the intercalative mode of interaction of all the complexes (1–18) with DNA. The shift in peak potential toward positive side could be attributed to the intercalation of the compounds into the stacked base pair pockets of DNA. The shift in E1/2 for complexes on binding to Cu2+–DNA suggests that both Cu(II) and Cu(I) forms bind to DNA but to different extents. Analogous to the treatment of the association of small molecules with micelles and DNA, the ratio of the equilibrium constants, K2+/K+ for the binding of the Cu(II) and Cu(I) forms of complexes to DNA can be estimated from the net shift in E1/2, assuming reversible electron transfer. For a Nernstian electron transfer in a system in which both the oxidized and reduced form associate with a

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third species such as DNA in solution, the following scheme can be applied. Here CuLn+–DNA represents the CuLn + complex bound to Cu2+–DNA.

From differential pulse voltammogram of the complexes (1 and 10) (Fig. 4), it is apparent that current intensity gets decreased upon the incremental addition of DNA. The shift in potentials both in presence and absence of DNA is related to the ratio of binding constant by the following equation: Eb −E f ¼ 0:0591 log K½red =K½oxd




where Eb and Ef are peak potentials of complex in presence (bound) and absence (free form) of DNA, respectively. The result of the present study is that all metal(II) complexes 1–18 show one electron transfer during the redox process. The ipa/ipc value of such redox process is less than unity which peculiarly supports the quasi-reversible nature of redox process which occurs on the surface of glassy carbon working

Fig. 5. Plot of relative specific viscosity vs. [complex]/[DNA] for Cu(II) complexes 1–3 and 10–12 in Tris–HCl buffer (pH 7.2) at 37 °C.

electrode. The synthesized complexes give both the anodic and cathodic peak potential shifts which are either positive or negative (Table S3). These shifts specify the intercalating mode of DNA binding with metal(II) complexes. 3.11. Viscometric Studies Hydrodynamic measurement, mainly viscosity, is a sensitive technique to determine the DNA binding mode. The binding nature of the complexes with DNA, viscosity measurements on the solutions of DNA incubated with the compounds have been carried out. The viscosity of CT-DNA solution increased after the addition of the solution of 1–18 and EB. With the [Complex]/[DNA] of complexes 1–3, 10–12 and EB increasing, the viscosity of the CT-DNA increased, as shown in Fig. 5. Intercalation of a species into DNA base pairs generally caused a significant increase in the viscosity of the DNA solution, due to an increase in the separation of the base pairs to accommodate the bound species, which was evidenced by a classical DNA intercalator like EB [35,36] and nonclassic intercalation under the same conditions typically caused either a less pronounced change (positive or negative) in DNA solution viscosity or none at all. The changes of the relative viscosity of CT-DNA bound to the complexes were similar to the known intercalator EB. The values of relative specific viscosities of DNA in the absence and presence of complexes 1 and 2 and the ligand are plotted against [complex]/ [DNA]. The increase of the relative viscosity of CT-DNA followed the order, 10 N 12 N 11 N 1 N 3 N 2 (Fig. 5). These results parallel the phenomena observed in the competitive binding studies. 3.12. Antimicrobial Assay

Fig. 4. Differential pulse voltammograms of complexes 1 (a) and 10 (b) in Tris–HCl buffer (pH 7.2) at 37 °C in presence of increasing amount of DNA (0–120 μM).

The synthesized Cu(II), Co(II), and Zn(II) complexes having cysteine and bpy/phen ligands have been evaluated for their antimicrobial actions. The minimum inhibitory concentration (MICs) and minimum bacterial/fungicidal concentration (MBC/MFC) values are given in Tables 1 and 2. The MIC is the lowest concentration of an antibacterial or antifungal compound that will inhibit the visible growth of microorganisms after period of incubation, and the minimum inhibitory concentrations are important in diagnostic laboratories to confirm the resistance of microorganisms to biologically active compounds. The experimental result indicates that all the complexes are having higher inhibition efficiency than their free ligands which can be explained on the basis of Overtone's concept and Chelation theory [37]. When the antimicrobial activity of metal complexes are investigated, the following

M. Selvaganapathy et al. / Journal of Photochemistry & Photobiology, B: Biology 157 (2016) 77–88 Table 1 MIC and MBC results of the test compounds (1–18) on three bacterial strains. Complex

1

L L2 L3 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Standard

Staphylococcus aureus

Escherichia coli

Bacillus subtilis

MIC

MBC

MIC

MBC

MIC

MBC

15.2 16.9 12.2 2.25 3.15 4.25 4.25 3.35 4.55 3.20 3.15 4.25 1.15 1.75 1.55 1.50 3.15 2.25 1.25 2.75 1.25 1.25

16.6 15.1 13.6 2.65 5.25 6.25 3.25 5.55 6.65 2.85 5.25 6.25 2.50 2.55 2.65 2.25 3.25 2.75 3.24 5.57 2.65 2.5

15.9 16.7 13.7 3.25 4.25 6.35 2.25 4.75 6.85 3.15 4.25 6.15 1.25 2.75 2.85 1.25 2.25 3.15 1.28 4.76 1.85 1.10

16.2 14.8 12.9 2.85 3.50 4.75 2.85 3.75 5.25 2.65 3.50 4.75 2.65 2.75 3.25 2.65 3.50 3.75 2.84 3.75 2.75 2.50

16.7 15.3 13.5 3.00 3.75 4.85 3.15 3.35 4.85 4.00 3.75 4.65 1.05 3.15 2.85 3.00 2.75 2.65 1.16 3.34 1.82 1.00

15.8 14.2 12.1 3.25 5.65 6.75 4.15 5.85 7.25 3.25 5.65 6.85 2.45 2.55 1.25 3.25 3.85 3.85 2.85 2.45 2.50 2.50

Amikacin is used as the standard. MIC (μM) = minimum inhibitory concentration, i.e. the lowest concentration to completely inhibit bacterial growth; MBC (μM) = minimum bactericidal concentration, i.e., the lowest concentration to completely kill bacteria.

factors [38,39] should be considered: (i) the chelate effect of ligands; (ii) the nature of the N-donar ligands; (iii) the total charge of the complex; (iv) existence and the nature of the ion neutralizing the ionic complex, and (v) the nuclearity of the metal centre in the complex. The tested complexes were more active against Gram-positive than Gram-negative bacteria (Table 1). It may be concluded that the antimicrobial activity of the compounds is related to cell wall structure of the bacteria. It is possible because the cell wall is essential to the survival of bacteria and some antibiotics are able to kill bacteria by inhibiting a step in the synthesis of peptidoglycan. Gram-positive bacteria possess a thick

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cell wall containing many layers of peptidoglycan and teichoic acids, but in contrast, Gram-negative bacteria have a relatively thin cell wall consisting of a few layers of peptidoglycan surrounded by a second lipid membrane containing lipopolysaccharides and lipoproteins. These differences in cell wall structure can produce differences in antibacterial susceptibility and some antibiotics can kill only Grampositive bacteria and is infective against Gram-negative pathogens. There are other factors which also increase the activity, which are solubility, presence of the bulkier organic moieties, conductivity, and bond length between the metal and the ligand. Current studies reveal that higher electronegativity and large atomic radius decrease the effective positive charges on the metal complex molecules and it results to higher antimicrobial activity [40]. The results of fungicidal screening (Table 2) show that metal chelates were highly active than the free ligands against phytopathogenic fungi. Furthermore, the mode of action of the compounds may involve the formation of a hydrogen bond through the azomethine nitrogen atom (N CN) with the active centers of cell constituents, resulting in interference with the normal cell process. The variation in the effectiveness of different compounds against different organisms depends either on the impermeability of the cells of the microbes or the difference in ribosomes of microbial cells [41]. These complexes also disturb the respiration process of the cell and thus block the synthesis of proteins, which restricts further growth of the organism. In general, metal complexes are more active than ligands as they may serve as major cytotoxic species and thus exhibiting their broad spectrum nature and can be further used in pharmaceutical industry for mankind, as an antimicrobial agent, after testing its toxicity to human beings. The observed antimicrobial activities of the synthesized complexes (Tables 1 and 2) are found to be higher than few of the analogue complexes found in the literature [42,43].

Table 2 MIC and MFC results of the test compounds (1–18) on three fungal strains. Complex

L1 L2 L3 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Standard

Aspergillus niger

Rhizoctonia bataticola

Candida albicans

MIC

MFC

MIC

MFC

MIC

MFC

16.2 16.5 12.8 3.30 3.26 4.62 3.73 3.57 4.85 4.30 3.65 4.68 2.25 2.57 2.85 3.10 3.26 2.58 2.32 3.57 2.35 2.25

17.6 14.8 13.1 2.62 4.38 5.05 3.07 4.75 5.65 3.62 4.85 5.25 2.55 2.75 2.65 3.12 3.38 2.75 3.07 3.75 2.65 2.50

16.5 15.7 13.4 3.02 3.74 4.38 3.43 4.06 4.72 4.02 3.74 4.32 2.03 2.06 2.72 3.02 3.54 2.36 2.43 2.06 2.72 2.00

16.8 15.1 12.9 3.65 3.27 4.75 3.75 3.56 5.25 3.65 3.25 4.25 2.25 2.56 3.25 2.65 3.27 2.75 2.75 2.56 3.25 2.65

16.2 15.7 13.5 2.86 2.95 3.83 4.64 3.92 4.02 4.36 2.85 3.83 1.25 1.42 2.02 1.36 2.55 2.83 2.64 2.92 2.02 1.25

15.8 14.2 12.5 2.53 4.15 5.25 3.85 4.72 5.65 3.53 4.10 5.25 2.75 2.52 2.65 2.53 3.15 3.25 2.85 2.72 2.65 2.50

Greseofulvin is used as the standard. MIC (μM) = minimum inhibitory concentration, i.e. the lowest concentration to completely inhibit fungal growth; MFC (μM) = minimum fungicidal concentration, i.e., the lowest concentration to completely kill fungus.

Fig. 6. a. The agarose gel electrophoresis of the Cu(II) complexes (10 μM) showing the cleavage of DNA in the presence of MPA (100 μM). Lane 1: DNA control (20 μM); Lane 2: DNA + L1 (100 μM); Lane 3: DNA + CuCl2 (500 μM), Lane 4: DNA + MPA (100 μM), Lanes 5–7: DNA + 10, 11, and 12 + MPA (100 μM); Lanes 8–10: DNA + 1, 2, and 3 + MPA (100 μM); Lane 11: DNA + EcoR1. b. The agarose gel electrophoresis of the Co(II) complexes (10 μM) showing the cleavage of DNA in the presence of MPA (100 μM). Lane 1: DNA control (20 μM); Lane 2: DNA + L2 (100 μM); Lane 3: DNA + CoCl2 (500 μM), Lanes 4–6: DNA + 4, 5, and 6 + MPA (100 μM); Lanes 7–9: DNA + 13, 14, and 15 + MPA (100 μM); Lane 10: DNA + EcoR1. c. The agarose gel electrophoresis of the Zn(II) complexes (10 μM) showing the cleavage of DNA in the presence of MPA (100 μM). Lane 1: DNA control (20 μM); Lane 2: DNA + L3 (100 μM); Lane 3: DNA + ZnCl2 (500 μM), Lanes 4–6: DNA + 16, 17, and 18 + MPA (100 μM); Lanes 7–9: DNA + 7, 8, and 9 + MPA (100 μM); Lane 10: DNA + EcoR1.

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Fig. 7a. Gel electrophoresis diagram showing the cleavage of supercoiled pBR322 DNA (20 μM) by 1, 2, and 3 with addition of external additives in DMF-Tris buffer medium and incubated at 37 °C for 1 h: Lane 1: DNA control, Lanes 2–4: DNA + 1, 2, and 3 + NaN3 (100 μM), Lanes 5–7: DNA + 1, 2, and 3 + SOD (1 U), Lanes 8–10: DNA + 1, 2, and 3 + DMSO (4 μL); Lane 11: DNA + EcoR1. b. Gel electrophoresis diagram showing the cleavage of supercoiled pBR322 DNA (20 μM) by 10, 11, and 12 with addition of external additives in DMF-Tris buffer medium and incubated at 37 °C for 1 h: Lane 1: DNA control, Lanes 2–4: DNA + 10, 11, and 12 + DMSO (4 μL), Lanes 5–7: DNA + 10, 11, and 12 + SOD (1 U), Lanes 8–10: DNA + 10, 11, and 12 + NaN3 (100 μM); Lane 11: DNA + EcoR1.

3.13. Chemical Nuclease Activity There has been considerable interest in DNA endonucleolytic cleavage activated by transition metal complexes. It is well known that DNA cleavage is reflected by relaxation of the supercoiled circular form (Form I) of pBR322 DNA resulting in nicked circular (Form II) and/or linear forms (Form III). In order to ascertain the nature of the DNA cleavage process, we have investigated its cleavage mechanism. The DNA cleavage activity of the complexes has been studied using a plasmid relaxation assay to monitor the conversion of circular supercoiled DNA (SC) to its nicked circular (NC) form. All complexes are inactive in the absence of any reducing agent when the reactions are carried out in the dark. As shown in Fig. 6a–c, the cleavage activity of the complexes 1–18 has been a crucial distinction in the presence of 3-mercaptopropionic acid (MPA), whereas the supercoil DNA (Form I) disappeared with the linear DNA (Form III) increasing noticeably. It indicated that the copper(II) complexes might be an admirable DNA cleavage agent in the presence of MPA as a reducing agent under the present experimental condition. The complexes (10 μM) show efficient DNA cleavage activity giving the order 10–18 (phen) N 1–9 (bpy). In order to explore the mechanistic pathway of the cleavage activity, comparative DNA cleavage experiment of complexes 1–18 were carried out in presence of some standard radical scavengers such as DMSO as hydroxyl radical scavenger (•OH), sodium azide (NaN3) as singlet

oxygen (1O2) quencher and superoxide dismutase (SOD) as superoxide anion radical O•− 2 scavenger, prior to the addition of complexes to DNA solution (Fig. 7a and b). The addition of DMSO (Fig. 7a and b) to Cu(II) complexes diminishes the cleavage activity which is indicative of the involvement of hydroxyl radical in the cleavage process. In the case of NaN3 and SOD, the Form II of plasmid DNA was converted to linear Form III indicating two subsequent and proximate single-strand breaks of DNA non-randomly. Similarly, Co(II) and Zn(II) complexes also showed inhibition of DNA cleavage in presence of DMSO, completely quenched the formation of band II (Fig. S6 and S7) suggesting the involvement of diffusible (HO) hydroxyl radicals as one of the ROS responsible for DNA breakage. On the other hand, addition of NaN3 and SOD did not show significant quenching of the cleavage revealing that singlet oxygen and superoxide anion were not involved in the cleavage process. In all the electrophoresis experiments, the last lane shows DNA digested with restriction enzyme EcoR1 which contains linear form of two DNA fragments. These cleavage patterns suggested their cleavage mechanism was possibly involving in hydroxyl radical species in the strand cleavage reactions [44]. 3.14. In Vitro Photochemosensitivity Cellular toxicity of the complexes 1–6 against four human cancer cell lines in dark and visible light was evaluated from the MTT assay. The complexes were generally found to show photocytotoxicity in these four cells in a visible light of 400–700 nm compared to the nonirradiated samples (Table 3). A dose-dependent anti-proliferative activity of the complexes was observed against all the cells in dark as expected due to reduction of copper(II) to copper(I) by thiols like GSH and cysteine and subsequent generation of radicals. DMSO is a biocompatible solvent used in several biological tests at lower concentration. The solvent (dimethyl sulphoxide, DMSO) shows no effect in cell growth. Photo-irradiation of the samples with visible light of 400–700 nm resulted in an enhanced cytotoxicity of the complexes. An enhancement in the cytotoxicity of the complexes in human cancer cells was observed when exposed to visible light compared to the samples in dark. Hence, a tenfold enhancement in cell cytotoxicity is observed upon photoexcitation in visible light. The phen complex 10 was most toxic to HeLa and MCF-7 cells with an IC50 value of 3.8 and 3.5 μM when exposed to visible light (Fig. 8). The IC50 values of the complexes along with Photofrin and cisplatin are given in Table 3 [45,46]. The observed photocytotoxicity of the complexes is comparable to that of Photofrin. Cisplatin lacking any photoactive moiety showed an IC50 value of ≈60–70 μM in both dark and light under similar experimental conditions. Photo exposure of the cells in absence of the complex showed no apparent reduction in the cell viability. A similar photo-enhanced anti-proliferative behavior of the complexes was seen on the HepG-2 and HEp-2 cancer cell lines

Table 3 IC50 (μM) values of test compounds (1–18) against various human cancer cell lines. Compounds

IC50 (μM) a

b

Dark

1 2 3 10 11 12 c Photofrin® Cisplatin a b c

Visible light

HeLa

HepG-2

MCF-7

HEp-2

HeLa

HepG-2

MCF-7

HEp-2

38.6 ± 0.4 39.8 ± 0.3 33.7 ± 0.4 27.4 ± 0.4 28.7 ± 0.6 31.5 ± 0.2 N 50 72.3 ± 0.4

39.3 ± 0.5 43.8 ± 0.6 39.1 ± 0.3 25.1 ± 0.6 31.7 ± 0.2 34.6 ± 0.4 N 50 63.6 ± 0.2

30.4 ± 0.1 33.2 ± 0.2 34.3 ± 0.6 21.2 ± 0.5 30.3 ± 0.2 34.1 ± 0.3 N 50 69.6 ± 0.5

32.3 ± 0.3 34.3 ± 0.5 36.8 ± 0.8 20.3 ± 0.3 27.2 ± 0.1 32.6 ± 0.2 N 50 57.4 ± 0.3

4.5 ± 0.1 4.1 ± 0.7 4.8 ± 0.4 3.8 ± 0.3 4.2 ± 0.6 4.3 ± 0.2 3.6 ± 0.3 68.7 ± 0.6

4.1 ± 0.3 6.5 ± 0.4 7.4 ± 0.6 4.7 ± 0.2 5.1 ± 0.6 8.2 ± 0.1 4.2 ± 0.4 58.2 ± 0.3

3.6 ± 0.3 5.5 ± 0.1 6.5 ± 0.2 3.5 ± 0.5 5.1 ± 0.2 6.1 ± 0.3 3.6 ± 0.7 64.1 ± 0.4

4.7 ± 0.2 4.5 ± 0.6 5.8 ± 0.3 4.1 ± 0.1 4.3 ± 0.2 5.1 ± 0.4 2.8 ± 0.5 53.5 ± 0.7

The IC50 values correspond to 24 h incubation in dark. The IC50 values correspond to 4 h incubation in dark followed by photo-exposure to visible light (400–700 nm, 10 J cm−2). The IC50 values (633 nm excitation; fluence rate: 5 J cm−2) of Photofrin® (converted to μM using the approximate molecular weight of Photofrin® = 600 g M−1).

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metal-based non-PDT test compound. Hence, it is clear that the noncovalently DNA binding complexes are excellent cytotoxic agents than the covalently DNA binding complexes. Hence, these copper(II) complexes which exhibit higher DNA binding affinity and prominent DNA cleavage activity display efficient cytotoxicity and anticancer properties. Further, these complexes are found to have better activity than few of similar type of complexes, reported in the literature [47,48]. The results showing remarkable photocytotoxicity of the phen copper(II) center are of significance in the virtually unexplored chemistry of bioorganometallics in PDT. 3.15. Lymphocyte Cytotoxicity Assay Complexes 1–6 were tested with a microlymphocytotoxicity assay for cell viability counting. The cytotoxic activity of complexes 1–6 against lymphocytes which are normal white blood cells isolated from human blood samples has been investigated. Interestingly, all the live healthy cells appear as green-colored fluorescence (Fig. 9) upon treatment with the complexes 1–6, which reveals that the complexes are non-toxic to normal cells (lymphocytes) as expected for a better drug. 4. Conclusion Fig. 8. Photo cytotoxicity of complex 10 (shown by circle) and phen ligand (shown by triangle) in human breast cancer MCF7 cells on 2 h incubation in dark ( ) followed by photo-irradiation for 45 min to visible light ( ) (400–700 nm) as determined by MTT assay.

on exposure to visible light. Again, the phen complex 10 was most cytotoxic toward HepG-2 and HEp-2 cells giving an IC50 of 4.1 and 4.7 μM in visible light. The activities of the complexes are similar in both the cell lines used. The ligands or the metal salts alone were significantly nontoxic both in dark and visible light. The bpy complexes 1–3 were found to be relatively non-toxic compared to other complexes and this could be due to their reduced uptake or quick efflux from the cancer cells. In the ranges of concentrations used, the obtained data indicate that the order of cytotoxic effect against four cell lines (Table 3) is 10 N 1 N 11 N 2 N 12 N 3, although the variable activity of the complexes may be due to oxidation–reduction potentials. Cisplatin is used as a

The synthesis and structural analysis of eighteen new Cu(II), Co(II), and Zn(II) complexes with 2,2′-bipyridine or 1,10-phenanthroline as a co-ligand is presented along with studies on their binding with DNA and their photocytotoxicity against tumor cell lines. An intercalative mode of binding between the octahedral complexes and CT-DNA is identified by UV–vis. (absorption and fluorescence emission measurements), cyclic voltammetry, differential pulse voltammetry, and viscometric measurements. Fascinatingly, the complexes 1–18 follow the hydrolytic cleavage pathway (hydroxyl radical (•OH)) with complex 10 displaying a higher efficiency than complex 1. In addition, the in vitro photocytotoxicity assay conducted on human cancer cell lines demonstrates that both the complexes are active against all the cancer cell lines. The phen complexes show significant PDT effect in visible light in MCF-7 cancer cells with slightly low cellular dark toxicity. The phen complex is significantly more photocytotoxic than the organic phen base. Copper being a bio-essential metal ion, its complexes showing significant photocytoxicity in visible light and comparatively low cellular dark toxicity could be of importance toward designing and developing new copper-based PDT agents. Moreover, all the complexes are very promising as drugs for potential clinical applications in cancer therapy as they damage DNA by double strand cleavage and cause cell death mainly through the apoptotic mode. Acknowledgments The authors express their sincere thanks to the UGC, New Delhi, for the financial assistance, and the managing board, principal, and head of the Department of Chemistry, VHNSN College, Virudhunagar, Tamil Nadu, for their constant encouragement and providing research facilities. Instrumental facilities provided by Sophisticated Analytical Instrument Facility (SAIF), IIT Bombay, and CDRI Lucknow are gratefully acknowledged. Appendix A. Supplementary data Supplementary data to this article can be found online at http://dx. doi.org/10.1016/j.jphotobiol.2016.02.008. References

Fig. 9. Phase contrast microscope image of live lymphocytes stained with acridine orange (AO) upon treatment with complex 10.

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