Syntheses, Characterization and Biological Activity of Coordination Compounds of 3-Hydroxy-2-methyl-4H-pyran-4-one and Its Mixed Ligand Complexes with 1,2-Diaminocyclohexane

Abstract

Coordination compounds of 3-hydroxy-2-methyl-4H-pyran-4-one with iron(III), cobat(III) and chromium(III) were synthesized with M:L (1:2). Mixed ligand coordination compounds of 3-hydroxy-2-methyl-4H-pyran-4-one and 1,2-diaminocyclohexane using the same metal ions were also synthesized M:L1:L2 (1:1:1) where L1 is 3-hydroxy-2-methyl-4H-pyran-4-one and L2 is 1,2-diaminocyclohexane. The coordination compounds obtained were characterized using electronic and infrared spectral analyses, magnetic susceptibility and percentage metal analysis. They were also evaluated for their cytotoxic and antioxidant activities. The result obtained suggested that octahedral geometry was obtained for all the compounds, as a result of additional two molecules of the solvent coordinated to the metal ions. Both the primary and secondary ligands coordinated in a bidentate fashion. The synthesized compounds exhibited moderate cytotoxicity, although none was as active as the standard. The cobalt(III) mixed ligand complex elicited the highest activity. The synthesized compounds all exhibited good to moderate antioxidant activity.

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Aiyelabola, T. (2021) Syntheses, Characterization and Biological Activity of Coordination Compounds of 3-Hydroxy-2-methyl-4H-pyran-4-one and Its Mixed Ligand Complexes with 1,2-Diaminocyclohexane. Advances in Biological Chemistry, 11, 106-125. doi: 10.4236/abc.2021.113008.

1. Introduction

The application of coordination compounds in chemistry and technology is many, varied and well established. Coordination compounds in the form of mixed ligand complexes have found usage commercially in medicine. These include cis-platin and auranofin used in anticancer chemotherapy and rheumatoid arthritis, respectively [1] [2] [3].

Cancer is a leading cause of death worldwide. It accounted for 8.2 million as deaths; around 22% of all deaths not related to communicable diseases. Deaths from cancer worldwide are projected to continue rising, with an estimated 13.1 million deaths in 2030 (about a 70% increase) [4] [5]. Chemotherapy with cytotoxic drugs serves as the main treatment modality for certain types of cancer [6] [7]. Studies have shown that the presence of increased level of exogenous antioxidants prevents certain types of free radical damage that has been associated with cancer development [8]. Previous data also suggested that certain antioxidants selectively inhibit the growth of tumor cells, induce cellular differentiation, and alter the intracellular redox state, thereby enhancing the effect of cytotoxic therapy [8] - [13]. Additionally, research has shown that antioxidants may also reduce certain types of toxicity associated with chemotherapeutic treatments [8] - [13]. However, it was inferred that they may do so by interfering with the efficacy of conventional therapy. Therefore, antioxidant supplementation during chemotherapy may compensate for treatment of cancer-induced antioxidant depletion, alleviate side effects and also maintain or improve general health and well-being [14]. A consequence of this is that drugs with antioxidant mechanisms are being widely proposed as starting point for the development of new therapeutic interventions in several pathological disorders associated with oxidative damage, caused by reactive oxygen species (ROS), such as cancer [15]. The search for metal-derived antioxidants has therefore received much attention in more recent times.

Phenolic compounds are a class of very important compounds with multiple biological functions including antioxidant activity which have been reported to be related to the radical scavenging ability of the hydroxyl group. A number of studies have reported the relative correlation between phenol and antioxidant activity [16] [17]. 3-Hydroxy-2-methyl-4H-pyran-4-one (L1), is a naturally occurring organic compound bearing a phenolic substituent, and it is used primarily as a flavor enhancer. Reports have shown that its derivatives exhibited limited in vitro antiproliferative activity towards cancer cells lines, with the mechanism of activity suggested as inducing apoptosis in the cancer cells [18] [19]. 1,2-Diaminocyclohexane (L2) is a compound that has found usage in the syntheses of anticancer drugs. It mixed ligand complexes are currently used commercially for the treatment of cancer and cisplatin-resistant tumours [20] [21]. These include tetraplatin [PtCl4(L2)] and oxaliplatin [Pt(L2)(oxalate)]. Earlier workers in an attempt to develop novel metal-based drugs with a different therapeutic profile to cis-platin, synthesized a series of tin compounds containing the L2, including the Sn analogue of tetraplatin. The result obtained suggested the complexes exhibited reasonable activity with increased differential toxicity across the cell line panel [22].

As a consequence it was considered to synthesize coordination compounds of 3-hydroxy-2-methyl-4H-pyran-4-one (L1) Figure 1, with iron(III), cobalt(III) and chromium(III). Additionally its mixed ligand complex with 1,2-diaminocyclohexane (L2), Figure 1, as the secondary ligand was synthesized. These were then characterized using electronic, infrared spectroscopy, magnetic susceptibility measurement and percentage metal analysis. The synthesized complexes were thereafter investigated for their cytotoxic activity using brine shrimp lethality assay, and antioxidant activity.

2. Materials and Method

2.1. General

All materials used are of high analytical grade. Melting points were determined in an open capillary tube on a Gallenkamp (Variable heater) melting point apparatus. The infrared spectra of all synthesized products and their ligands were obtained using Agilent Cary 630 FTIR. Magnetic susceptibility measurement of the metal complexes was using a MSB Mk1 magnetic susceptibility balance, Sherwood Scientific and corrected with. The electronic spectra, of all the compounds, were obtained in solution, in the wavelength range 400 - 1000 nm using 1800 Shimadzu ultra-violet spectrophotometer. The metal analyses for all synthesized compounds were obtained using titrimetric method using EDTA. The cytotoxic analysis of the compounds was carried out using brine shrimp lethality assay. The antioxidant activity for all the complexes was carried out using four assays namely; nitric oxide radical inhibition, ferric reducing antioxidant power (FRAP), ferrous ion-chelating ability and cupric ion reducing antioxidant capacity assays. The compounds were synthesized using an adaptation of previous method [23]. The equations of the reactions are given in Equations (1)-(6).

FeCl 2 + 2 ( L 1 ) + H 2 O 2 [ Fe ( L 1 ) 2 ( H 2 O ) 2 ] (Compound 1) (1)

FeCl 2 + L 1 + L 2 + H 2 O 2 [ Fe ( L 1 ) ( L 2 ) ( H 2 O ) 2 ] (Compound 2) (2)

CoCl 2 + 2 L 1 + H 2 O 2 [ Co ( L 1 ) 2 ( H 2 O ) 2 ] (Compound 3) (3)

CoCl 2 + L 1 + L 2 + H 2 O 2 [ Co ( L 1 ) ( L 2 ) ( H 2 O ) 2 ] (Compound 4) (4)

Figure 1. (a) L1 = 3-hydroxy-2-methyl-4H-pyran-4-one; (b) L2 = 1,2-diaminocyclohexane.

CrCl 3 + 2 L 1 [ Cr ( L 1 ) 2 ( H 2 O ) 2 ] (Compound 5) (5)

CrCl 3 + L 1 + L 2 [ Cr ( L 1 ) ( L 2 ) ( H 2 O ) 2 ] (Compound 6) (6)

where: L1 = 3-Hydroxy-2-methyl-4H-pyran-4-one; L2 = 1,2-diaminocyclohexane.

2.2. Syntheses of Compounds

1) Synthesis of Compound 1: An aqueous solution of Iron(II) chloride hexahydrate (1.59 g, 0.01 M) was poured into a flat bottom flack and was heated with stirring. To this was added, drop-wise, hydrogen peroxide (0.36 g, 0.01 M). Aqueous ethanolic solution of 2-hydroxy-2-methyl-4H-pyran-4-one (2.36 g, 0.02 M) was added drop-wise. The reaction mixture was heated for 2 hr, during which a pale brown precipitate was obtained. This was filtered washed first with methanol then diethyl ether and dried in a desiccator. Yield: 2.58 g (75.36%), M.pt/d.t: 135˚C - 136˚C (d), metal composition (%.): 15.73 (found); 16.03 (calcd). The complex was soluble in water but sparingly soluble in ethanol, methanol. Similar procedure was used for the preparation of the under listed complexes.

2) Synthesis of Compound 2: Iron(II) chloride hexahydrate (1.58 g, 0.01 M) in a flat bottom flask to which was added hydrogen peroxide (0.35 g, 0.01 M), 2-hydroxy-2-methyl-4H-pyran-4-one (1.18 g; 0.01 M) and 1,2-diaminocyclohexane (0.80 g, 0.01 M) heated with stirring. This gave a dark green precipitate. Yield: 2.29 g (69.25%), M.pt/d.t: 221˚C - 222˚C, metal composition (%.): 17.02 (found); 16.61 (calcd). The complex was soluble in water but sparingly soluble in ethanol, methanol.

3) Synthesis of Compound 3: Cobalt(II) chloride hexahydrate (1.61 g, 0.01 M) to which, hydrogen peroxide (0.36 g, 0.01 M) and 2-hydroxy-2-methyl-4H-pyran-4-one (2.38 g; 0.02 M) was added and heated with stirring. This gave a light pink precipitate. Yield: 2.49 g (71.89%), M.pt/d.t: 182˚C - 183˚C, metal composition (%.): 17.23 (found); 17.00 (calcd.). The complex was soluble in water but sparingly soluble in ethanol, methanol.

4) Synthesis of Compound 4: Cobalt(II) chloride hexahydrate (1.63 g, 0.01 M), hydrogen peroxide (0.34 g, 0.01 M) and 2-hydroxy-2-methyl-4H-pyran-4-one (1.18 g; 0.01 M) and 1,2-diaminocyclohexane (1.18 g; 0.01 M) gave a pink precipitate. Yield: 2.42 g (73.31%), M.pt/d.t: 112˚C - 113˚C (d), metal composition (%.): 17.44 (found); 17.61 (calcd.). The complex was soluble in ethanol, methanol but sparingly soluble in water.

5) Synthesis of Compound 5: Chromium(III) chloride hexahydrate (1.58 g, 0.01 M) in a flat bottom flask to which was added and 2-hydroxy-2-methyl-4H-pyran-4-one (2.38 g; 0.02 M) heated with stirring. This gave a dark green precipitate. Yield: 2.33 g (68.52%), M.pt/d.t: 164˚C - 165˚C (d), metal composition (%.): 15.56 (found); 15.29 (calcd.). The complex was soluble in water but sparingly soluble in ethanol, methanol.

6) Synthesis of Compound 6: Chromium(III) chloride hexahydrate (1.58 g, 0.01 M) in a flat bottom flask to which was added 2-hydroxy-2-methyl-4H-pyran-4-one (1.18 g; 0.01) and 1,2-diaminocyclohexane (1.80 g, 0.01 M) heated with stirring. This gave a dark green precipitate. Yield: 1.72 g (62.20%), M.pt/d.t: 127˚C - 128˚C (d), metal composition (%.): 15.80 (found); 15.85 (calcd.). The complex was soluble in ethanol, methanol but sparingly soluble in water.

2.3. Cytotoxicity Bioassay

The procedure used was modified from the assay described by Solis et al. [24]. Brine shrimp (Artemia salina) were hatched from shrimp eggs in a conical shaped vessel (1 L). Subsequently the vessel was filled with sterile, artificial seawater under continuous aeration for 48 h. After hatching, active nauplii free from eggshells were collected from brighter portion of the hatching chamber. These were employed for the assay. Ten nauplii were drawn through a Pasteur pipette and placed in each vial containing 4.5 mg/L of brine solution. In each experiment, different volumes of the sample chelates were added to 4.5 mL of brine solution. This produced different concentrations of 20, 40, 60, 80 and 100 µg/mL. Solutions were maintained at room temperature for 24 h under light. The surviving larvae were counted. Experiments were conducted along with the control (vehicle treated), of the test substances in a set of three tubes per dose. Estimation of the LC50 values was estimated using probit® analysis on a USEPA computer program.

2.4. Antioxidant Activity

The antioxidant activity for all the compounds was carried out using four assays.

2.4.1. Inhibition of Nitric Oxide Radical

The inhibition nitric oxide radical activity of the synthesized compounds was carried out according to the method of Green et al., (1982) as described by Marcocci et al. (1994) [25]. The reaction mixture, containing 0.1 ml of different concentrations (10, 5, 2.5, 1.25, 0.625, 0.3125 mg/ml) of the sample and 0.9 ml of sodium nitroprusside (2.5 mM) in phosphate buffer saline was incubated under illumination for 150 minutes. After incubation, 0.5 ml of 1% sulphanilamide in 5% phosphoric acid was added and incubated in the dark for 10 min., followed by addition of 0.5 ml 0.1% NED (N-1-napthyl ethylenediamine dihydrochloride). The absorbance of the chromophore formed was measured at 546 nm (Marcocci et al., 1994) [25]. The percentage inhibition of nitric oxide radical formation was calculated as expressed below.

% inhibition = [ ( A control A sample ) / A control ] × 100

where Acontrol = absorbance of control sample and Asample = absorbance of a tested sample.

2.4.2. Ferrous Ion-Chelating Ability Assay

The ferrous ion-chelating (FIC) assay was carried out according to the method of Singh and Rajini, 2004 with some modifications [26]. Solutions of 2 mM FeCl2·4H2O and 5 mM ferrozine were diluted 20 times. An aliquot (1 ml) of different concentrations of sample was mixed with 1ml FeCl2·4H2O. After 5 min incubation, the reaction was initiated by the addition of ferrozine (1 ml). The mixture was shaken vigorously and after a further 10 min incubation period the absorbance of the solution was measured spectrophotometrically at 562 nm. The percentage inhibition of ferrozine–Fe2+ complex formations was calculated by using the formula:

Chelating effect % = [ ( A control A sample ) / A control ] × 100

where Acontrol = absorbance of control sample (the control contains FeCl2 and ferrozine, complex formation molecules) and Asample = absorbance of a tested sample.

2.4.3. Determination of Ferric Reducing Antioxidant Power (FRAP)

The FRAP assay uses antioxidants as reductants in a redox-linked colorimetric method with absorbance measured with a spectrophotometer (Benzie and Strain, 1999) [27]. A 300 mmol/L acetate buffer of pH 3.6, 10 mmol/L 2,4,6-tri-(2-pyridyl)-1,3,5-triazine and 20 mmol/L FeCl3·6H2O were mixed together in the ratio of 10:1:1 respectively, to give the working FRAP reagent. A 50 μl aliquot of the sample at 0.1 mg/ml and 50 μl of standard solutions of ascorbic acid (20, 40, 60, 80, 100 µg/ml) was added to 1 ml of FRAP reagent. Absorbance measurement was taken at 593 nm exactly 10 minutes after mixing against reagent blank containing 50 µl of distilled water.

All measurements were taken at room temperature with samples protected from direct sunlight. The reducing power was expressed as equivalent concentration (EC) which is defined as the concentration of antioxidant that gave a ferric reducing ability equivalent to that of the ascorbic acid standard.

2.4.4. CUPRAC Assay

In order to determine the cupric ions (Cu2+) reducing ability of compounds, the method of Apak et al. was used with little modification as described by Gulcin [28] [29]. 0.25 ml CuCl2 solution (0.01 M), 0.25 ml ethanolic neocuproine solution (7.5 × 103 M), and 0.25 ml CH3COOH4 buffer (1 M) were added to a test tube, followed by mixing with 0.25 ml of extracts. The total reaction volume was adjusted to 2 ml with distilled water, and the solution was mixed well. The tubes were stoppered and kept at room temperature for 30 min, and absorbance was measured at 450 nm. Increased absorbance indicates increased reduction capability which is express as Trolox equivalent (TEAC) using Trolox as standard.

3. Results and Discussion

3.1. FTIR

3.1.1. FTIR of the Ligands

1) The FTIR spectrum of the primary ligand L1, exhibited a sharp band at 3330 cm−1. This is indicative of the O-H, stretching frequency of the phenolic substituent of the ligand. Supporting this were bands observed at 2100 and 1250 cm−1, which were assigned to ν(C-OH) and δ(O-H)Ar, respectively [30] [31]. The spectrum also showed bands at 1650 and 1530 cm−1 were attributed to ν(C=O) +ν(C=C). This is suggestive of conjugation between the aldehyde moiety of L1 and the aromatic C=C. Bands observed at 1630 and 1450 cm−1 were ascribed to ν(C=O) + δ(C-H). The band observed at 1230 cm−1 was assigned to ν(C-O-C), from the ether substituent of the ligand [30] [31].

2) Ligand L2: The spectrum of the secondary ligand L2, showed a broad band at 3340 cm−1. This is attributable to the ν(N-H). Two medium bands observed at 1590 and 1420 cm−1 were ascribed to δ(N-H) [30] [31]. A weak band at 1180 cm−1 was assigned to ν(C-N) [30] [31].

3.1.2. FTIR of the Complexes

1) Compound 1: On coordination to the metal ion, the O-H stretching frequency of the phenolic substituent of the ligand shifted to higher frequency by 28 cm−1 at 3358 cm−1. Another weak band was also observed at 3242 cm−1 (Figure 2). Both bands suggested possible intermolecular hydrogen bonding [30] [31] [32]. Two bands observed at 3768 and 3432 cm−1 were assigned to hydroxyl stretching frequency, indicative of probable intermolecular hydrogen bonding. The bands attributed to ν(C=O) + ν(C=C) shifted to lower and higher frequencies at 1613 and 1558 cm−1 in this complex [30] [31] [32]. This therefore served as evidence for the coordination of the oxygen atom of the aldehyde moiety to the metal ion. In support of this was the shift observed for the ν(C=O) + δ(C-H) to 1457 cm−1 [30] [31] [32]. Another band observed at 1394 cm−1 was attributed δ(O-H)Ar, of the phenol substituent of the ligand. A shift of 52 and 32

Figure 2. FTIR spectrum of compound 1.

cm−1 was observed for the ν(C-O-C). This may be attributable to the elongation of the bonds in this particular substituent as a result of coordination of substituents adjacent to it. This suggests the rearrangement of the ring as a result of coordination. In addition to the bands observed in the free base ligand, two bands were observed at 1073 and 1021 cm−1 and were assigned to δ(O-H). Metal-oxygen frequency band was observed at 685 cm−1, which is an additional support of coordinationof the ligand to the metal ion [32].

2) Compound 2: The spectrum of compound 2 elicited two broad frequency bands at 3779 and 3570 cm−1 (Figure 3). The latter frequency band was attributed to the phenolic O-H, which shifted to higher frequency by 270 cm−1, in comparison with that obtained for ligand L1. Therefore, this served as indication of the coordination of the oxygen atom of the phenol substituent to the metal ion [32]. This was supported by the δ(O-H)Ar which shifted to 1397 cm−1 in the spectrum of compound 2. On the other hand, the former ν(O-H) frequency band suggested the presence and coordination of an hydroxyl substituent in addition to the phenolic OH of the primary ligand [30]. This was supported by the band observed at 1080 cm−1 assignable to δ(O-H) [30] [31] [32]. On coordination the ν(C=O) + ν(C=C) frequency band shifted to 1651 and 1574 cm−1. The ν(C=O) + δ(C-H) shifted to 1606 and 1457 cm−1. Evidence of the coordination of the secondary ligand L2 to the metal ion was given by the frequency bands observed at 3410 and 1181 cm−1, which were attributed to ν(N-H) and ν(C-N). This was further supported by a band observed in the region ~510 cm−1. Additional support for the formation of a mixed ligand complex, coordination of the primary ligand L1 and secondary ligand L2 and the solvent was suggested by the four bands observed in the region ~600 cm−1. These were attributed to metal-oxygen and metal-nitrogen stretching frequency (Table 1) [32].

Figure 3. FTIR spectrum of compound 2.

Table 1. Relevant infrared spectra bands for the ligands and complexes (cm−1).

3) Compound 3: This spectrum exhibited two bands at 3358 and 3246 cm−1, which are ascribable to ν(O-H) (Figure 4). Similar to what was obtained for compound 2, the higher frequency band was ascribed to the phenolic ν(O-H) which is supported by the appearance of the δ(O-H)Ar at 1394 cm−1. The lower frequency band was attributed to an additional ν(O-H) from the solvent. Bands observed at 1073 and 1021 cm−1 due to δ(O-H), served as evidence in support of this. Coordination of the carbonyl oxygen atom to the central metal ion was suggested by the shifts in the (C=O) + ν(C=C) at 1613 and 1558 cm−1 and the ν(C=O) + δ(C-H) at 1457 cm−1. Bands observed at 685, 551 and 514 cm−1, attributable toν(M-O), suggested coordination to the central metal ion via three different atoms namely: two oxygen atoms from the carbonyl and phenol substituent of L1 and oxygen atoms from the solvent [32].

4) Compound 4: The ν(O-H) was observed at 3205 and 3104 cm−1, assignable to the phenolic substituent of L1 and hydroxyl substituent from the solvent. Evidence in support of this was provided by frequency bands observed at 1364 and 1069, 1013 cm−1 (Figure 5), which corresponds to δ(O-H)Ar and δ(O-H) respectively. This therefore, suggested the coordination of donor atoms from this substituent to the central metal ion. Shifts were observed for the bands ν(C=O) + ν(C=C) to 1661, 1555 cm−1 and ν(C=O) + δ(C-H) to 1605, 1457 cm−1. As such, indicating coordination of the ligand to the metal ion. Evidence for the coordination of the secondary ligand and the formation of a mixed ligand complex was provided by the band at 1182 cm−1 assigned to δ (N-H) + ν(C-N). Bands observed at 685, and 520 cm−1 were ascribed to ν(M-O) and ν(M-N) respectively. This suggested the coordination of both oxygen and nitrogen atoms to the central metal ion [32].

5) Compound 5: In the spectrum of compound 5 bands observed at 3459 and 3153 cm−1 were attributed to phenolic substituent of L1 and hydroxyl substituent from the solvent. The δ(O-H)Ar was observed at 1397 cm−1 and δ(O-H) at 1066

Figure 4. FTIR spectrum of compound 3.

Figure 5. FTIR spectrum of compound 4.

and 1013 cm−1. In support of this was the ν(C-OH) observed at 2117 cm−1 (Figure 6). The ν(C=O) + ν(C=C) was observed at 1558 cm−1 while the ν(C=O) + δ(C-H) was observed at 1606 and 1457 cm−1. Coordination of the metal ion to the ligand was confirmed by the bands observed at 685 and 568 cm−1 [32].

6) Compound 6: Evidence obtained from the spectrum of compound 6 suggested the coordination of the phenolic substituent of L1 and hydroxyl substituent from the solvent (Table 1). In addition to this the formation of a mixed ligand complex was suggested by the appearance of double spike at 3116 and 3063 cm−1 (Figure 7). This may be attributed to the nitrogen-hydrogen stretching

Figure 6. FTIR spectrum of compound 5.

Figure 7. FTIR spectrum of compound 6.

frequency, from the –NH2 moiety. This was corroborated by the presence of medium bands at 1494 and 1092 cm−1 assignable to δ(N-H) and ν(C-N) frequency bands respectively. Bands observed at 629 and 536 cm−1 were ascribed to ν(M-O) and ν(M-N) respectively [32].

3.2. Electronic Spectra

1) Compound 1 and 2: The electronic spectra of compounds 1 and 2 exhibited three intense bands in the ultraviolet region respectively (Table 2). These were assigned to n → σ*, n → π* and π → π* for compound 1 and n → σ*, n → π* and π → π* for compound 2 [30] [31]. The visible region revealed weak broad frequency bands 437 and 798 nm for compound 1. These were assigned to 6A1g4T1g and 6A1g4Eg. For compound 2 the visible region exhibited bands at 626 and 717 nm ascribable to 6A1g4T2g and 6A1g4Eg. Iron(III) ion in an octahedral environment has a d5 configuration with a spin only magnetic moment of 5.92 BM [33]. Compounds 1 and 2 exhibited a magnetic moment of 6.83 and 6.45 BM, respectively. Both are consistent with an octahedral geometry and therefore corroborates the results obtained for their electronic spectra [33] [34].

2) Compound 3 and 4: The spectrum for compound 3 elicited an intense band at 275 nm (Table 2) assigned to n → π* transition. The spectrum for compound 4 however exhibited bands at 224, 274, 307 and 326 nm corresponding to n → σ*, n → π* and π → π*, π → π* transitions [30] [31]. The visible spectrum for compound 3 showed frequency bands at 444, 607 and 760 nm which were given the assignment 1A1g1T2g, 1A1g1T1g and 1A1g1Eg. On the other hand the spectrum of compound 4 showed bands at 824 and 878 nm and ascribed to 1A1g1Eg and 1A1g1A2g d-d transitions [33] [34]. Cobalt(III) ion in an octahedral environment has a d6 configuration, with a spin only magnetic moment of 4.84 BM. Compounds 3 and 4 elicited magnetic moment of 5.22 and 5.82 BM respectively. This is in agreement with previous reports and the high values obtained in comparison with the spin only values may be attributed to orbital contribution [33] [34].

3) Compound 5 and 6: The electronic spectrum of compound 5 showed four bands in the visible region at 416, 592,760 and 803 nm assigned to 4A2g(F) → 4T1g(F), 4A2g(F) → 4T2g(F), 4A2g(F) → 4T1g(P) and 4A2g(F) → 1Eg d-d transitions respectively [33]. On the other hand Compound 6 elicited two broad at 430 and 601 nm, which were attributed to 4A2g(F) → 4T1g(P) and 4A2g(F) → 4T1g(F) transitions. The ultraviolet region exhibited three intense bands at 220, 270 and 360 nm for compound 6 [30] [31]. These were ascribed to n → σ*, n → π* and π → π* transitions [30] [31]. Compound 5 on the other hand elicited two intense bands in the ultraviolet region at 269 and 215 nm, attributable to n → π* and π → π* transitions [30] [31]. Magnetic moment of 3.76 and 3.33 BM was obtained for compounds 5 and 6 respectively. Both values are consistent with octahedral geometry. Although this is lower than the spin only magnetic moment for chromium(III) ion in an octahedral environment, 3.87 BM, this may be as a result of antiferromagnetism [33] [34].

Based on the results obtained it is suggested that L1 coordinated in a bidentate fashion coordinating to the metal ion via the oxygen atoms of the phenol moiety and the carbonyl substituent. On the other hand L2 coordinated using both nitrogen atoms of its amino substituents. Typically, ν(O-H) ought to exhibit two bands in the high energy region, indicating symmetric and asymmetric stretch. Two bands were observed within this range for this vibrational frequency. However, the presence of both (O-H)Ar and δ(O-H) suggested the coordination of the oxygen atom from the hydroxyl substituent of the solvent. This was further supported by the results obtained from the electronic spectra. The result obtained

Table 2. Relevant electronic spectra bands (nm), for the ligands and complexes.

suggested octahedral geometry for all the complexes. Although compounds 1, 3 and 5 were designed as binary complexes, with M:L1 of (1:2). However, ternary complexes were obtained, with additional water molecules coordinated to the central metal ions, to attain octahedral geometry (Figure 8). For the mixed ligand complexes ternary complexes were designed with M:L1:L2, 1:1:1. However, in this case, coordination of the primary and secondary ligands occurred as well as coordination of two molecules of water to give an octahedral geometry (Figure 9). These were suggested from the infrared and electronic spectra of the compounds in addition to the magnetic moment and percentage metal concentration obtained for the compounds. The obtained percentage metal composition was in good agreement with the calculated values. The octahedral geometry assumed by the compounds served as an indication of the stability of the octahedral geometry for the central metal ions. Iron(III), cobalt(III) and chromium(III) have d5, d6 and d3 configurations respectively, electron distribution will therefore favour partially-, singly- or fully-filled sub-shells, as a result of stability. Therefore, generally, with few exceptions, these metal ions in their +3 oxidation states assume the octahedral geometry.

3.3. Cytotoxicity

Although there are many methods used in the study of anticancer activity, brine shrimp lethality assay serves as a preliminary assay for the cytotoxici abilities for probable potent compounds [35] [36]. Brine shrimp lethality assay of the synthesized compounds and ligands was carried out, and the result obtained indicated that both ligands, the synthesized compounds and the metal salts were less potent in comparison with the standard K2Cr2O7 (LC50 5.56 μg/ml). However, both ligands exhibited significant (𝑃 < 0.05) cytotoxic activity with L1 and L2 exhibiting LC50 of 14.42 and 13.39 μg/ml respectively. This indicated L2 was the more active of the ligands and is in good agreement with previous studies. The chloride salt of iron(III), cobalt(III) and chromium(III) also elicited reasonable cytotoxic activity with LC50 9.73, 12.88 and 12.62 μg/ml respectively. The order of activity for the compounds was 4 > 1 > 6 > 3 > 2 > 5 with LC50 33.53, 36.15,

Figure 8. Structural representation for compound 1, 3 and 5 (M-Fe(III), Co(III) and Cr(III) and S = H2O).

Figure 9. Structural representation for compound 2, 4 and 6 (M-Fe(III), Co(III) and Cr(III) and S = H2O).

38.31, 41.89, 53.25 and 64.51 μg/ml. The mixed ligand complexes exhibited better active in comparison to their binary ligand counterparts. However compound 2, was an exception to this. In this case the ternary complex compound 1 was more cytotoxic than the mixed ligand complex, compound 2. Since the result from the structural characterization indicated that all the compounds assumed similar geometry, this difference may be ascribed to the size of the Fe(III). Generally it may be inferred that chelation reduced the toxicity of the ligand and metal ion [37].

3.4. Antioxidant Activity

It is widely recognized that many of today’s diseases are due to the oxidative stress that results from an imbalance between formation of ROS/RNS and their neutralization when endogenous antioxidant mechanisms are unable to quench the free radicals [8] [9] [10]. The free radicals are known to be scavenged by antioxidants; therefore, the search for effective antioxidants has become crucial. However, the determination of the total antioxidant activity of a compound is a complex procedure with varying chemical diversity [8] [38]. This usually occurs through several mechanisms which are influenced by many factors, which cannot be fully described with one single method. Therefore, it is essential to perform more than one type of antioxidant capacity measurement to take into account the various mechanisms of antioxidant action [38] [39] [40] [41]. In this study, four complementary tests were used to assess the antioxidant activity of compounds 1-6 viz.: inhibition of nitric oxide radical, determination of ferric reducing antioxidant power (FRAP), ferrous ion-chelating ability and cupric ion reducing antioxidant capacity assays.

1) Inhibition of nitric oxide radical: Nitric oxide is a diffusible free radical which acts as an effector molecule in a large number of biological processes [8] [42]. Conversely, excessive production and release of NO is reported to be linked to cytotoxicity. NO is generated in biological tissues by specific nitric oxide synthase which metabolize arginine to citrulline with the formation of NO via a five electron oxidative reaction. NO has been shown to be directly scavenged by phenolic compounds [8] [42] [43]. In this assay nitric oxide, generated from sodium nitroprusside in aqueous solution at physiological pH, interacts with oxygen to produce nitrite ions which was measured by Griess reaction. Therefore, this is a hydrogen atom transfer (HAT)-based assay. Compounds 1-6 exhibited good NO scavenging activity leading to the reduction of the nitrate concentration in the assay medium. The order of potency is given by 2 > 4 > 1 > 6 > 5 > 3. The NO scavenging capacity was concentration dependent with 1.00 mg/ml scavenging most efficiently. Compound 2 significantly (P < 0.05) inhibited the accumulation of nitrite (Table 3) and was the most comparable to the standard, ascorbic acid with 61.95 µg/ml.

2) Ferrous Ion-chelating Ability: Iron chelating activity is one of the major characteristics of antioxidants [13]. Elevated levels of iron have been implicated in changing hydrogen peroxide ion, which is less reactive to hydroxyl radical, a highly toxic free radical and thereby inducing oxidative stress [8] [43]. In view of the aforementioned effects of iron overload, the manufacture of drugs to combat this challenge would be a promising strategy to ameliorate the consequences of oxidative stress and as well reduce the concentration of transition metals as catalysts of lipid peroxidation [8] [43]. From the results, all the synthesized compounds generally exhibited above average activity (>50%). However, the compounds were not as potent as EDTA, the standard. Although compounds 3 and 4 demonstrated good chelating abilities eliciting as much as 73.77% and 68.21% inhibition respectively at 1 mg/ml comparable with the standard (Table 4). The order of activity was observed as 3 > 4 > 6 > 2 > 1 > 5.

3) Determination of ferric reducing antioxidant power FRAP: The principle of this method is based on the reduction of a colourless ferric-tripyridyltriazine

Table 3. Nitric oxide radical inhibition ability of the compounds.

Table 4. Ferrous ion-chelating ability of the compounds.

Table 5. Ferric reducing antioxidant power (FRAP) of the compounds based on ascorbic acid equivalent.

Table 6. Cupric ions reducing ability of compounds expressed as trolox equivalent.

complex to its blue ferrous coloured form owing to the action of electron donation in the presence of antioxidants [8] [28]. The compounds were observed as the ferric-tripyridyltriazine complex in a dose-dependent manner. Compound 4 elicited the highest activity of 1.03 mg/g. The order at which the complex was reduced by the compounds was observed as 4 > 6 > 5 > 1 > 2 > 3 (Table 5).

4) Cupric ion reducing antioxidant capacity (CUPRAC): In a comprehensive review by Prior et al. the authors classified CUPRAC as one of the electron-transfer based methods, and also summarize the superiorities of the CUPRAC method over other antioxidant assays [44]. The chromogenic oxidizing reagent bis(neocuproine)copper(II) chloride (Cu(II)-Nc), reacts with the antioxidants agent. In this reaction, the reactive Ar-OH groups of polyphenolic antioxidants are oxidized to the corresponding quinones (ascorbic acid is oxidized to dehydroascorbic acid) [44]. In this reaction compound 4 exhibited the highest potency with a Trolox equivalent of 0.86 (Table 6). All the compounds exhibited generally significant activity (P <0.05) with order of reactivity observed as 4 > 3 > 2 > 6 > 1 > 5.

From the results obtained for both biological analyses, compound 4 elicited the best cytotoxic activity. Additionally, it exhibited the best antioxidant activity for two of the four essays and the second best for the other two. It may therefore be inferred from these results that a good cytotoxic agent may also have antioxidant activity. Further studies are however suggested to ascertain their complementary effect.

4. Conclusion

The result obtained demonstrated the relative stability of the octahedral geometry for iron(III), cobalt(III) and chromium(III). It was suggested that two molecules of the solvent coordinated to the metal ions in addition to the ligands to attain octahedral geometry. Both the primary and secondary ligands coordinated in a bidentate fashion. The primary ligand 3-Hydroxy-2-methyl-4H-pyran-4-one coordinated using the oxygen atom of its phenolic substituent and the oxygen atom of the carbonyl, to act as an OO'ligand. The results obtained also suggested that the phenolic substituent was not deprotonated prior to coordination. Therefore L1 acted as a neutral ligand. The secondary ligand 1,2-diaminocyclohexane also coordinated using both nitrogen atoms of its amino substituent. The synthesized compounds exhibited moderate cytotoxicity, although they were not as active as the standard. The cobalt(III) mixed ligand complex elicited the highest activity and also antioxidant activity.

Conflicts of Interest

The authors declare no conflicts of interest regarding the publication of this paper.

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