# Synthesis, Characterization, Cyclic Voltammetry, and Biological Studies of Co(II), Ni(II), and Cu(II) Complexes of a Tridentate Schiff Base, 1-((E)-(2-Mercaptophenylimino) Methyl) Naphthalen-2-ol (H2L1)

Authors: Maurice Kuate; Mariam Asseng Conde; Evans Ngandung Mainsah; Awawou G. Paboudam; Francis Merlin M. Tchieno; Kevin I. Y. Ketchemen; Ignas Tonle Kenfack; Peter T. Ndifon
Journal: Journal of Chemistry
 (2020)

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## Abstract

A novel tridentate Schiff base, 1-((E)-(2-mercaptophenylimino) methyl) naphthalen-2-ol (H2L1), was synthesized by the condensation reaction of 2-hydroxy-1-naphthaldehyde with 2-aminothiophenol in absolute ethanol. The resulting ligand was reacted with Co(II), Ni(II), and Cu(II) ions to obtain tetrahedral CoL1, NiL1, and square planar CuL1 complexes. The Schiff base and its metal complexes were characterized using 1H-NMR, microanalysis, FT-IR, UV-visible, and mass spectroscopy (ESI-MS). All the compounds are soluble in DMSO and DMF. Spectroscopic studies show that the ligand coordinates to the metal center through the azomethine nitrogen, naphthoxide oxygen, and thiophenoxide sulfur to form a tridentate chelate system. Conductance measurements show that these compounds are molecular in solution. Cyclic voltammetry studies show Co(III)/Co(II) and Cu(II)/Cu(I) redox systems to be quasi-reversible involving a monoelectronic transfer while Ni(III)/Ni(II) was irreversible. In vitro antibacterial and antifungal activity against five bacterial strains (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, Enterococcus faecalis, and Proteus mirabilis) and five fungal strains (Candida albicans, Candida glabrata, Candida tropicalis, Candida krusei, and Candida parapsilosis) showed no antifungal activity but moderate antibacterial activity on E. coli, S. aureus, P. aeruginosa, and P. mirabilis bacterial strains. Antioxidant studies reveal that the ligand and its Cu(II) complex are more potent than Co(II) and Ni(II) complexes to eliminate free radicals.


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## Body

## 1. Introduction

The synthesis, structure, and reactivity of Schiff bases have generated a lot of interest due to their ease of preparation and structural flexibility. Schiff bases are known to possess potential applications in biological modeling, catalysis, design of molecular magnets, and the synthesis of therapeutic agents [1–4]. Schiff bases have also been shown to exhibit a broad range of biological activities, including antifungal, antibacterial, antioxidant, and anti-inflammatory properties [4–9]. The common structural feature of Schiff’s bases is the presence of the characteristic azomethine functionality, RHC = NR′, which has been shown to account for the observed pharmacological activities [4]. Schiff bases of salicylaldehyde derivative behave generally as bidentate ligands and coordinate through the oxygen atom of the phenolic group and the nitrogen atom of the azomethine group. Schiff bases are found to be the most appropriate chelating ligands in coordination chemistry due to their ability to bond through N, O, and/or S atoms [5, 6] as either bidentate or tridentate chelators which typically form four- or six-coordinate complexes.The stabilization of metal centers using bidentate mixed donor Schiff bases involving N- andO-donors has been extensively studied [5–8]. Most tridentate donor ligands have been shown to stabilize metal ion centers, forming strain-free five- or six-membered rings [9, 10].The increasing resistance of microbes to antibacterial and antifungal drugs has necessitated the search for new compounds to target these pathogens. The incorporation of metal-based systems into organic molecules is expected to enhance the biological properties of these drugs. Complexes of Schiff bases have been extensively studied as potential antibacterial, antifungal, and anticancer agents [11–13]. Tridentate Schiff base complexes having oxygen, nitrogen, and sulfur donor atoms have been reported to possess unusual structure [14], antimicrobial [15, 16], anti-inflammatory [17, 18], antiviral, anticancer [19], and electrochemical properties [20, 21]. In recent years, tridentate Schiff bases derived from aminophenol received enormous attention as ligands due to their metal-binding ability and their various applications. Our group has recently embarked on studies on the biological activities of complexes of heterocyclic Schiff base ligands [15, 22]. One such class of heterocyclic Schiff base ligands is the pyridine carboxaldehyde isonicotinoyl hydrazones, derived from isoniazid (isonicotinic acid hydrazide) which are generally bidentate or tridentate chelators and typically form six-coordinate complexes [22–24].In continuation of our studies on metal complexes of aromatic Schiff base ligands, we report, herein, the synthesis and electrochemical studies of Co(II), Ni(II), and Cu(II) complexes of a tridentate Schiff base derived from 2-aminothiophenol and 2-hydroxy-1-naphthaldehyde and the evaluation their antioxidant and antimicrobial properties.


## 2. Experimental

### 2.1. Materials and Measurements

All reagents and solvents were obtained from commercial sources and used without any further purification. Microanalysis (C, H, and N) data were obtained using a Perkin-Elmer model 240C elemental analyzer. Infrared spectra were determined using KBr pellets on an FTIR spectrophotometer Shimadzu model 8400s in the region 4000–400 cm−1. Electronic spectra were recorded in DMSO on UV-Vis spectrophotometer spectroScan 80D in a 200–800 nm range. 1H-NMR for the ligand was recorded on a Bruker AMX 300 spectrometer operating at 300.13 MHz using DMSO-d6 as solvent and TMS as a standard reference. The high-resolution mass spectrum was obtained with a Waters Micromass LCT Premier Mass Spectrometer in Electron Spin ionization (ESI) mode. The cyclic voltammograms were recorded on a μ-Autolab III potentiostat using a conventional three-electrode cell. A glassy carbon electrode was used as the working electrode while Ag/AgCl and platinum gauze were employed as the reference and auxiliary electrodes, respectively. A 0.1 M H2SO4 solution was used as the supporting electrolyte, and all synthesized compounds were dissolved in DMF. All electrochemical measurements were carried out in solutions containing 10−3 M for either the ligand or its metal complexes. These solutions were degassed with N2 prior to any recording of cyclic voltammograms while maintaining them under an inert atmosphere.


### 2.2. Synthesis

The Schiff base ligand and its complexes were synthesized according to the general synthetic procedure [25, 26].


#### 2.2.1. Synthesis of Schiff Base (H2L1)

The Schiff base (H2L1) was synthesized by the condensation of 2-aminothiophenol and 2-hydroxy-1-naphthaldehyde. An ethanolic solution of 2-hydroxy-1-naphthaldehyde (3,271 g; 19 mmol) was slowly added to an ethanolic solution of 2-aminothiophenol (2 mL; 19 mmol). The reaction mixture was maintained under reflux for two hours. The volume of the resulting yellow solution was reduced by evaporation and cooled in ice water. The yellow precipitate formed was filtered, washed several times with ethanol, and air-dried. The yellow product was collected: Yield: 92%; m.p 160°C; Analysis for C17H13NOS (%); Found (calc): C, 73.00 (73.09); H, 4.43 (4.69); N, 4.90 (5.01); ESI-MS in methanol: m/z 279.2.


#### 2.2.2. Synthesis of the Complexes

All complexes were synthesized using the same procedure. 1.00 mmol of metal nitrate hydrate was dissolved in 10 mL ethanol and added dropwise to a vigorously stirring (1.00 mmol) methanolic solution of ligand H2L1. The resulting solution was stirred under reflux for two hours and cooled. The colored precipitate of each complex was obtained, filtered, washed thoroughly with ethanol, and air-dried at room temperature.


### 2.3. Electrochemical Studies

Cyclic voltammetry studies were carried out on aμ-Autolab (type III) potentiostat, with a three-electrode cell containing a glassy carbon as the working electrode, an Ag/AgCl reference electrode, and Pt foil counter electrode. Prior to each experiment, the glassy carbon electrode was properly cleaned by polishing using an Al2O3 suspension. Solutions of the Schiff base ligand and complexes (1.0 × 10−3 M) in DMF with sulfuric acid (0.1 M) as a supporting electrolyte were degassed using dry nitrogen gas for 15 minutes and then blanketed with the same gas during the experiments. All compounds were investigated at room temperature. The voltammograms were recorded at a potential scan rate of 100 mV·s−1.


### 2.4.In Vitro Antioxidant Activity

The ligand, H2L1, and its complexes were tested for in vitro antioxidant activities at 37°C using DPPH free radical scavenging assay [27] with some modifications. Solutions of the ligand and complexes at different concentrations (200, 100, 50, 25.5, and 12.25 μg/mL) were prepared in methanol as solvent. 1 mL of each sample at different concentrations and 540 μL of DPPH (0.08 mg/mL) solution were measured into different test tubes, and the mixture was shaken vigorously for about 2-3 minutes. The contents of the test tubes were then incubated in the dark for 30 minutes at room temperature. A blank DPPH solution without the sample used for the baseline correction gave a strong absorption maximum at 517 nm (purple color with ε = 8.32 × 103 M−1·cm−1). After incubation, the absorbance value for each sample 510–520 nm was measured using a UV-visible spectrometer. The observed decrease in absorbance values indicates that the compounds show scavenging activity and the relative free radical scavenging effects were calculated using the formula [28]:(1)scavenging effects%=Abscontrol−AbssampleAbscontrol×100,where Abs(control) is the absorbance of DPPH radical + DMF and Abs(sample) is the absorbance of DPPH radical + sample (test samples/standard). Each analysis was made in triplicate, and the obtained mean data were compared with Trolox values used as standard.


### 2.5. Antimicrobial Study

The Schiff base ligand and its respective metal(II) complexes were screened forin vitro antibacterial and antifungal activity against some bacterial strains (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, Enterococcus faecalis, and Proteus mirabilis) and some fungal strains (Candida albicans, Candida glabrata, Candida tropicalis, Candida krusei, and Candida parapsilosis); Ciprofloxacin and Ketoconazole (at 1 mg/ml in DMSO) were used as antibacterial and antifungal references, respectively.


#### 2.5.1. Determination of the Diameters of Zone of Inhibition

The diameters of the zone of inhibition of the synthesized ligand and its complexes were determined using the agar well diffusion method [29, 30]. The stock solutions (1 mg/mL) of the compounds were prepared in DMSO. For the determination of the zone of inhibition, the medium was poured into the Petri dish and allowed to solidify at room temperature. Wells were made on the solidified medium, and the prepared solutions were added to the wells and allowed to diffuse into the wells. The indicator organisms were overlaid on to the agar medium, and the plates were incubated at 37°C for 24 hours for the bacteria and 48 hours for the yeasts. The antimicrobial activities were assessed by measuring the diameter of the zone of inhibition of the bacterial and fungal growth around every well with a ruler following two axes. For every product tested, three determinations were done.


#### 2.5.2. Minimum Inhibitory Concentration

The minimum inhibitory concentrations (MICs) were determined by the method of microdilution in a liquid environment in the 96-well microtiter plates for the bacterial species [31]. The quantitative antifungal activities of the test compounds were evaluated using the microdilution broth method [32]. Twofold serial dilutions of the compounds were prepared in 96-well microtiter plates using sterile nutrient broth as diluent. The plates were inoculated with 100 μL bacterial or fungal suspensions containing 1.5 × 108 colony-forming units (CFUs) [32] and incubated at 37°C for 48 hours for fungal species and 24 hours for bacterial species. The MIC value was defined as the lowest concentration of the compounds giving complete inhibition of visible growth.


## 2.1. Materials and Measurements

All reagents and solvents were obtained from commercial sources and used without any further purification. Microanalysis (C, H, and N) data were obtained using a Perkin-Elmer model 240C elemental analyzer. Infrared spectra were determined using KBr pellets on an FTIR spectrophotometer Shimadzu model 8400s in the region 4000–400 cm−1. Electronic spectra were recorded in DMSO on UV-Vis spectrophotometer spectroScan 80D in a 200–800 nm range. 1H-NMR for the ligand was recorded on a Bruker AMX 300 spectrometer operating at 300.13 MHz using DMSO-d6 as solvent and TMS as a standard reference. The high-resolution mass spectrum was obtained with a Waters Micromass LCT Premier Mass Spectrometer in Electron Spin ionization (ESI) mode. The cyclic voltammograms were recorded on a μ-Autolab III potentiostat using a conventional three-electrode cell. A glassy carbon electrode was used as the working electrode while Ag/AgCl and platinum gauze were employed as the reference and auxiliary electrodes, respectively. A 0.1 M H2SO4 solution was used as the supporting electrolyte, and all synthesized compounds were dissolved in DMF. All electrochemical measurements were carried out in solutions containing 10−3 M for either the ligand or its metal complexes. These solutions were degassed with N2 prior to any recording of cyclic voltammograms while maintaining them under an inert atmosphere.


## 2.2. Synthesis

The Schiff base ligand and its complexes were synthesized according to the general synthetic procedure [25, 26].


### 2.2.1. Synthesis of Schiff Base (H2L1)

The Schiff base (H2L1) was synthesized by the condensation of 2-aminothiophenol and 2-hydroxy-1-naphthaldehyde. An ethanolic solution of 2-hydroxy-1-naphthaldehyde (3,271 g; 19 mmol) was slowly added to an ethanolic solution of 2-aminothiophenol (2 mL; 19 mmol). The reaction mixture was maintained under reflux for two hours. The volume of the resulting yellow solution was reduced by evaporation and cooled in ice water. The yellow precipitate formed was filtered, washed several times with ethanol, and air-dried. The yellow product was collected: Yield: 92%; m.p 160°C; Analysis for C17H13NOS (%); Found (calc): C, 73.00 (73.09); H, 4.43 (4.69); N, 4.90 (5.01); ESI-MS in methanol: m/z 279.2.


### 2.2.2. Synthesis of the Complexes

All complexes were synthesized using the same procedure. 1.00 mmol of metal nitrate hydrate was dissolved in 10 mL ethanol and added dropwise to a vigorously stirring (1.00 mmol) methanolic solution of ligand H2L1. The resulting solution was stirred under reflux for two hours and cooled. The colored precipitate of each complex was obtained, filtered, washed thoroughly with ethanol, and air-dried at room temperature.


## 2.2.1. Synthesis of Schiff Base (H2L1)

The Schiff base (H2L1) was synthesized by the condensation of 2-aminothiophenol and 2-hydroxy-1-naphthaldehyde. An ethanolic solution of 2-hydroxy-1-naphthaldehyde (3,271 g; 19 mmol) was slowly added to an ethanolic solution of 2-aminothiophenol (2 mL; 19 mmol). The reaction mixture was maintained under reflux for two hours. The volume of the resulting yellow solution was reduced by evaporation and cooled in ice water. The yellow precipitate formed was filtered, washed several times with ethanol, and air-dried. The yellow product was collected: Yield: 92%; m.p 160°C; Analysis for C17H13NOS (%); Found (calc): C, 73.00 (73.09); H, 4.43 (4.69); N, 4.90 (5.01); ESI-MS in methanol: m/z 279.2.


## 2.2.2. Synthesis of the Complexes

All complexes were synthesized using the same procedure. 1.00 mmol of metal nitrate hydrate was dissolved in 10 mL ethanol and added dropwise to a vigorously stirring (1.00 mmol) methanolic solution of ligand H2L1. The resulting solution was stirred under reflux for two hours and cooled. The colored precipitate of each complex was obtained, filtered, washed thoroughly with ethanol, and air-dried at room temperature.


## 2.3. Electrochemical Studies

Cyclic voltammetry studies were carried out on aμ-Autolab (type III) potentiostat, with a three-electrode cell containing a glassy carbon as the working electrode, an Ag/AgCl reference electrode, and Pt foil counter electrode. Prior to each experiment, the glassy carbon electrode was properly cleaned by polishing using an Al2O3 suspension. Solutions of the Schiff base ligand and complexes (1.0 × 10−3 M) in DMF with sulfuric acid (0.1 M) as a supporting electrolyte were degassed using dry nitrogen gas for 15 minutes and then blanketed with the same gas during the experiments. All compounds were investigated at room temperature. The voltammograms were recorded at a potential scan rate of 100 mV·s−1.


## 2.4.In Vitro Antioxidant Activity

The ligand, H2L1, and its complexes were tested for in vitro antioxidant activities at 37°C using DPPH free radical scavenging assay [27] with some modifications. Solutions of the ligand and complexes at different concentrations (200, 100, 50, 25.5, and 12.25 μg/mL) were prepared in methanol as solvent. 1 mL of each sample at different concentrations and 540 μL of DPPH (0.08 mg/mL) solution were measured into different test tubes, and the mixture was shaken vigorously for about 2-3 minutes. The contents of the test tubes were then incubated in the dark for 30 minutes at room temperature. A blank DPPH solution without the sample used for the baseline correction gave a strong absorption maximum at 517 nm (purple color with ε = 8.32 × 103 M−1·cm−1). After incubation, the absorbance value for each sample 510–520 nm was measured using a UV-visible spectrometer. The observed decrease in absorbance values indicates that the compounds show scavenging activity and the relative free radical scavenging effects were calculated using the formula [28]:(1)scavenging effects%=Abscontrol−AbssampleAbscontrol×100,where Abs(control) is the absorbance of DPPH radical + DMF and Abs(sample) is the absorbance of DPPH radical + sample (test samples/standard). Each analysis was made in triplicate, and the obtained mean data were compared with Trolox values used as standard.


## 2.5. Antimicrobial Study

The Schiff base ligand and its respective metal(II) complexes were screened forin vitro antibacterial and antifungal activity against some bacterial strains (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, Enterococcus faecalis, and Proteus mirabilis) and some fungal strains (Candida albicans, Candida glabrata, Candida tropicalis, Candida krusei, and Candida parapsilosis); Ciprofloxacin and Ketoconazole (at 1 mg/ml in DMSO) were used as antibacterial and antifungal references, respectively.


### 2.5.1. Determination of the Diameters of Zone of Inhibition

The diameters of the zone of inhibition of the synthesized ligand and its complexes were determined using the agar well diffusion method [29, 30]. The stock solutions (1 mg/mL) of the compounds were prepared in DMSO. For the determination of the zone of inhibition, the medium was poured into the Petri dish and allowed to solidify at room temperature. Wells were made on the solidified medium, and the prepared solutions were added to the wells and allowed to diffuse into the wells. The indicator organisms were overlaid on to the agar medium, and the plates were incubated at 37°C for 24 hours for the bacteria and 48 hours for the yeasts. The antimicrobial activities were assessed by measuring the diameter of the zone of inhibition of the bacterial and fungal growth around every well with a ruler following two axes. For every product tested, three determinations were done.


### 2.5.2. Minimum Inhibitory Concentration

The minimum inhibitory concentrations (MICs) were determined by the method of microdilution in a liquid environment in the 96-well microtiter plates for the bacterial species [31]. The quantitative antifungal activities of the test compounds were evaluated using the microdilution broth method [32]. Twofold serial dilutions of the compounds were prepared in 96-well microtiter plates using sterile nutrient broth as diluent. The plates were inoculated with 100 μL bacterial or fungal suspensions containing 1.5 × 108 colony-forming units (CFUs) [32] and incubated at 37°C for 48 hours for fungal species and 24 hours for bacterial species. The MIC value was defined as the lowest concentration of the compounds giving complete inhibition of visible growth.


## 2.5.1. Determination of the Diameters of Zone of Inhibition

The diameters of the zone of inhibition of the synthesized ligand and its complexes were determined using the agar well diffusion method [29, 30]. The stock solutions (1 mg/mL) of the compounds were prepared in DMSO. For the determination of the zone of inhibition, the medium was poured into the Petri dish and allowed to solidify at room temperature. Wells were made on the solidified medium, and the prepared solutions were added to the wells and allowed to diffuse into the wells. The indicator organisms were overlaid on to the agar medium, and the plates were incubated at 37°C for 24 hours for the bacteria and 48 hours for the yeasts. The antimicrobial activities were assessed by measuring the diameter of the zone of inhibition of the bacterial and fungal growth around every well with a ruler following two axes. For every product tested, three determinations were done.


## 2.5.2. Minimum Inhibitory Concentration

The minimum inhibitory concentrations (MICs) were determined by the method of microdilution in a liquid environment in the 96-well microtiter plates for the bacterial species [31]. The quantitative antifungal activities of the test compounds were evaluated using the microdilution broth method [32]. Twofold serial dilutions of the compounds were prepared in 96-well microtiter plates using sterile nutrient broth as diluent. The plates were inoculated with 100 μL bacterial or fungal suspensions containing 1.5 × 108 colony-forming units (CFUs) [32] and incubated at 37°C for 48 hours for fungal species and 24 hours for bacterial species. The MIC value was defined as the lowest concentration of the compounds giving complete inhibition of visible growth.


## 3. Results and Discussion

The ligand H2L1 was prepared by condensing equimolar amounts of 2-aminothiophenol with 2-hydroxy-1-naphthaldehyde in absolute ethanol under reflux (see Scheme 1).Scheme 1
Reaction equation of synthesis of the Schiff base.The corresponding Co(II), Ni(II), and Cu(II) complexes were synthesized in 1 : 1 molar ratio of the Schiff base ligand and the metal(II) ions using M(NO3)2 × H2O. The ligand is soluble in methanol and hot ethanol while both the ligand and its complexes are soluble in DMF and DMSO. The analysis and physical properties of the prepared compounds are given in Table 1.Table 1
Analysis and physical properties of Schiff base and the corresponding complexes.
CompoundsColorYield (%)Melting point (°C)Molar mass (g·mol−1)% found (calculated)Λm (Ω−1·cm2·mol−1)CHNH2L1Yellow92.00160279.3673.00 (73.09)4.43 (4.69)4.90 (5.01)0.6CuL1Dark green61.70˃360681.7761.34 (59.90)3.03 (3.25)4.07 (4.11)2.0CoL1Dark brown65.50˃360686.0659.47 (59.52)3.42 (3.45)3.96 (4.08)1.2NiL1Red64.40˃360672.0760.64 (60.76)3.28 (3.30)3.98 (4.17)0.8The yellow color of the ligand has been used to predict the formation of the Schiff base. Elemental analysis values for the ligand and its complexes (Table1) tallied with the expected values, thus confirming the purity of the compounds and the formation of the Schiff base ligand. A 1 : 1 reaction ratio for metal ion and the ligand during the formation of the complexes was proposed. Thus, the complexes were of the form [ML], which may thus dimerize [33, 34], to form a neutral dimeric four-coordinate complex [ML]2. The dimerization is presumed to occur as a result of bridging between the two M(II) centers through the deprotonated thiophenolic sulfur atoms. The very low molar conductivities of the synthesized compounds in DMSO indicated that all complexes are molecular in nature.


### 3.1. Infrared Study

The infrared spectrum of the ligand is presented in Figure1 and the data are summarized in Table 2. The infrared spectrum of the ligand exhibits absorption bands around 3255 and 2943 cm−1 which are attributed to the hydroxyl and thiol groups, respectively. The broad nature of the bands indicates the existence of an intramolecular hydrogen bonding between the O-H and nitrogen atom of -C=N groups [34]. The loss of absorption band at 3317–3385 cm−1 for ν(N-H) stretch of 2-aminothiophenol and at 1706 cm−1 for ν(C=O) of 2-hydroxy-1-naphthaldehyde and appearance of a peak at 1612 cm−1 attributed to azomethine group ν(C=N) in the ligand are indicative of the formation of the Schiff base, H2L1.Figure 1
Infrared spectrum of the ligand H2L1.Table 2
Infrared data of the Schiff base (H2L) and its corresponding complexes.
CompoundsνC=NνC-OνC-SνC-S asymνO-HνS-HνM-NνM-OνM-SνH2OH2L11612123873881732552943————CuL116001348741824——522561456—CoL116001360747837——5015554683452NiL116081375746821——508575454—The IR spectrum of the ligand has been compared with the IR spectra of metal complexes (Figure S1) that revealed the binding modes of the ligand to metal ions, which is confirmed by the change in the positions of absorption bands. The strong band at 1612 cm−1 attributed to the imine ν(C=N) functional group of the free ligands was redshifted, to 1608–1600 cm−1, in the complexes and indicative of the involvement of the imine nitrogen atom in coordination [34, 35]. The disappearance of the absorption bands around 3255 and 2943 cm−1 attributed to O-H and S-H groups in the spectra of the complexes is due to deprotonation and involvement of the oxygen and sulfur atoms in the coordination to metal ion [34, 35], thus confirming the tridentate nature of the ligand. This is further supported by the appearance of three new bands at 522–501 cm−1, 575–555 cm−1, and 468–454 cm−1 in the far-infrared spectra of the complexes, assigned to ν(M-N), ν(M-O) and ν(M-S), respectively [34–36]. In addition, the broad absorption peak observed at 3452 cm−1 in the Co(II) complex is attributed to O-H of the water molecule of crystallization.


### 3.2.1H NMR Spectral Analysis

1H-NMR spectra of the ligand and metal complexes confirmed the binding of the Schiff base with metal ions (Ni(II), Co(II), and Cu(II)). The intensities of each signal in the 1H-NMR spectra of the ligand and its metal complexes were proportional as expected to the number of different types of protons present. The 1H-NMR data of the compounds and the chemical shifts of the different types of protons are listed in Table 3.Table 3
1H-NMR data of ligand and its complexes.
Compoundsδ(O-H)δ(N=C-H)δ(S-H)δ(Ar-H)δ(H2O)H2L115.01 ppm (s, 1H)9.78 ppm (s, 1H)3.35 ppm (s, 1H)7.89–7.02 (m, 10H)—CuL1—8.09 ppm (s, 1H)—7.25–6.45 (m, 10H)—CoL1—8.36 ppm (s, 1H)—7.59–6.69 (m, 10H)3.34 ppm (s, 2H)NiL1—8.31 ppm (s, 1H)—7.78–6.69 (m, 10H)—The1H-NMR spectrum of the ligand is presented in Figure 2. The phenolic proton of the ligand appears downfield as a singlet at 15.01 ppm while the signal at 3.35 ppm was attributed to the thiol (-SH) proton. The broadness of the signal was due to a strong hydrogen bonding between the imine nitrogen and the phenolic proton [34–36]. On the other hand, the azomethine proton (-HC=N-) appeared as a strong singlet at 9.78 ppm, confirming the formation of the Schiff base. All the aromatic protons were observed at 7.89–7.02 ppm [38].Figure 2
1H-NMR spectrum of the Schiff base ligand H2L1.A comparison of1H-NMR spectra of copper, cobalt, and nickel complexes with that of the free ligand displayed an upfield shift in the signal of azomethine proton which appeared at 8.09–8.36 ppm suggesting the coordination of azomethine nitrogen to the metal ions [37]. However, signals at 15.01 ppm and 3.35 ppm corresponding to a hydroxyl group and thiol group, respectively, in free ligand were found to be absent in both complexes, indicating the deprotonation of the hydroxyl group of naphthaldehyde moiety and thiol group of 2-aminothiophenol [38]. However, the peak in the range 10-11 ppm corresponds to the possible enamine tautomerization proton. Moreover, the signal observed at 3.34 ppm, corresponding to two protons in Co(II) complex spectrum, is assigned to one water molecule of crystallization [38]. The above conclusion is in agreement with the IR analysis which confirmed that the ligand is tridentate and coordinates to the metal using the azomethine nitrogen, phenolic oxygen, and thiophenolic sulfur atoms. The 1H-NMR data of the ligand and its complexes are presented in Table 3.


### 3.3. Mass Spectra

The ESI-MS of the Schiff base was performed to determine the molecular weight and fragmentation pattern. The molecular ion peak was observed atm/z 279.2 confirming formula weight (C17H13NOS) of H2L1. The peak at m/z 302.2 can be assigned to [M + Na]+ and the peak at m/z = 304.1 is assigned to [MH + Na]+ probable (13C-isotope in this case, due to its very low relative abundance) [39]. The mass spectrum of H2L1 is shown in Figure 3.Figure 3
Mass spectra of the Schiff base ligand H2L1.


### 3.4. Electronic Spectra

The UV-visible spectra of compounds are presented in Figure4, and their electronic data in DMSO are given in Table 4. The spectrum of H2L1 exhibited high-intensity bands at 362 nm (27624 cm−1), assigned to π ⟶ π∗ transition, and at 465 nm (21505 cm−1) assigned to n ⟶ π∗ transition [39, 40]. The band at 362 nm in the spectrum of the Schiff base experienced a blue shift in the spectra of the complexes. The spectrum of CuL1 complex shows an absorption band at 586 nm attributed to 2B1g ⟶ 2A1g transition of square planar geometry around Cu(II) while the band observed at 436 nm was attributed to charge transfer transition [40]. The spectrum of CoL1 complex shows an absorption band at 640 nm attributed to 4A2 ⟶ 4T1(P) transition which corresponds to tetrahedral geometry, while the band observed at 476 nm corresponds to charge transfer transition [41–44]. The spectrum of NiL1 showed absorption bands at 564 nm, attributed to 3T2(F) ⟶ 3T1(F) transition which corresponds to tetrahedral geometry, and the band observed at 453 nm is charge transfer transitions [39].Figure 4
Electronic spectrum of ligand H2L1 and its corresponding complexes.
(a)(b)(c)(d)Table 4
Electronic spectral data of the ligand and its corresponding complexes.
Compoundsλ (nm) andν (cm−1)ε (L·mol−1·cm−1)AssignmentsSuggested structuresH2L1362 nm (27624 cm−1)2901π ⟶ π∗—465 nm (21505 cm−1)740n ⟶ π∗CuL1332 nm (30120 cm−1)1610H2L1 chromophore (π ⟶ π∗)Square planar436 nm (22935 cm−1)1113Charge transfer (L ⟶ M)586 nm (17064 cm−1)402B1g ⟶ 2A1gCoL1339 nm (29498 cm−1)1137H2L1 chromophore (π ⟶ π∗)Tetrahedral476 nm (21598 cm−1)578Charge transfer (L ⟶ M)640 nm (15625 cm−1)1764A2 ⟶ 4T1 (P)NiL1329 nm (30395 cm−1)626H2L1 chromophore (π ⟶ π∗)Tetrahedral453 nm (22075 cm−1)772Charge transfer (L ⟶ M)564 nm (17730 cm−1)463T1(F) ⟶ 3T2 (F)


### 3.5. Thermal Analysis

The differential and thermogravimetric analyses of CuL1, CoL1, and NiL1 complexes were determined under a nitrogen gas inert atmosphere in the range 0 and 600°C. The thermograms and data obtained are shown in Figure 5 and Table 5, respectively. The thermal decomposition curves of CuL1 and NiL1 complexes showed single step decomposition corresponding to the loss of two molecules of ligand at 385°C (found: 72.11; calc: 72.56) and 428°C (found: 68.36; calc: 68.55), respectively, with a DTA of 400°C and 430°C, respectively. The thermal decomposition curve of the CoL1 complex showed two-step decomposition. The first step at 220°C gave a loss (found: 15.12; calc: 15.11) which corresponds to one molecule of aminothiophenol, while the second step occurred in the range of 400–556°C (found: 44.05; calc: 44.10) corresponding to the loss of one molecule of aminothiophenol and two molecules of 2-hydroxy-1-naphthaldehyde [39].Figure 5
TGA and DTA of CuL1, NiL1, and CoL1 complexes.
(a)(b)(c)Table 5
Thermoanalytical results for CuL1, CoL1, and NiL1 complexes.
CompoundsStepTGA temperature (°C)DTA temperature (°C)Weight loss (%)ProcessFragment lostFoundCalcCuL1Single385400 (endothermic)72.1172.56DecompositionC17H13NO (2 moles)Residue>60023.0323.34Cu2SCoL11220384 (endothermic)15.1215.11DecompositionC6H7NS (1 mole) + 3/4H2O2400–556457 (endothermic)44.0544.10DecompositionC11H8O2 (2 moles) +C6H7NS (1 mole)Residue>600—40,1241.52Co2O3; CoS; CONiL1Single428430 (endothermic)68.3668.55DecompositionC17H13NS (2 moles)Residue>60030.9130.96Ni3O2


### 3.6. Proposed Structures of Complexes

In agreement with microanalysis and all the spectroscopic results, the tridentate Schiff base coordinated to the metal using nitrogen atom of the azomethine group (-HC=N-), oxygen atom of hydroxynaphthaldehyde, and sulfur atom of the aminothiophenol. During the coordination, complexes probably dimerized to form neutral dimeric four-coordinate systems [33]. The dimerization is presumed to occur as a result of the bridging between the M(II) centers through the deprotonated thiophenolic sulfur atoms as described in Figure 6.Figure 6
Proposed structure of the M(II) complexes.


## 3.1. Infrared Study

The infrared spectrum of the ligand is presented in Figure1 and the data are summarized in Table 2. The infrared spectrum of the ligand exhibits absorption bands around 3255 and 2943 cm−1 which are attributed to the hydroxyl and thiol groups, respectively. The broad nature of the bands indicates the existence of an intramolecular hydrogen bonding between the O-H and nitrogen atom of -C=N groups [34]. The loss of absorption band at 3317–3385 cm−1 for ν(N-H) stretch of 2-aminothiophenol and at 1706 cm−1 for ν(C=O) of 2-hydroxy-1-naphthaldehyde and appearance of a peak at 1612 cm−1 attributed to azomethine group ν(C=N) in the ligand are indicative of the formation of the Schiff base, H2L1.Figure 1
Infrared spectrum of the ligand H2L1.Table 2
Infrared data of the Schiff base (H2L) and its corresponding complexes.
CompoundsνC=NνC-OνC-SνC-S asymνO-HνS-HνM-NνM-OνM-SνH2OH2L11612123873881732552943————CuL116001348741824——522561456—CoL116001360747837——5015554683452NiL116081375746821——508575454—The IR spectrum of the ligand has been compared with the IR spectra of metal complexes (Figure S1) that revealed the binding modes of the ligand to metal ions, which is confirmed by the change in the positions of absorption bands. The strong band at 1612 cm−1 attributed to the imine ν(C=N) functional group of the free ligands was redshifted, to 1608–1600 cm−1, in the complexes and indicative of the involvement of the imine nitrogen atom in coordination [34, 35]. The disappearance of the absorption bands around 3255 and 2943 cm−1 attributed to O-H and S-H groups in the spectra of the complexes is due to deprotonation and involvement of the oxygen and sulfur atoms in the coordination to metal ion [34, 35], thus confirming the tridentate nature of the ligand. This is further supported by the appearance of three new bands at 522–501 cm−1, 575–555 cm−1, and 468–454 cm−1 in the far-infrared spectra of the complexes, assigned to ν(M-N), ν(M-O) and ν(M-S), respectively [34–36]. In addition, the broad absorption peak observed at 3452 cm−1 in the Co(II) complex is attributed to O-H of the water molecule of crystallization.


## 3.2.1H NMR Spectral Analysis

1H-NMR spectra of the ligand and metal complexes confirmed the binding of the Schiff base with metal ions (Ni(II), Co(II), and Cu(II)). The intensities of each signal in the 1H-NMR spectra of the ligand and its metal complexes were proportional as expected to the number of different types of protons present. The 1H-NMR data of the compounds and the chemical shifts of the different types of protons are listed in Table 3.Table 3
1H-NMR data of ligand and its complexes.
Compoundsδ(O-H)δ(N=C-H)δ(S-H)δ(Ar-H)δ(H2O)H2L115.01 ppm (s, 1H)9.78 ppm (s, 1H)3.35 ppm (s, 1H)7.89–7.02 (m, 10H)—CuL1—8.09 ppm (s, 1H)—7.25–6.45 (m, 10H)—CoL1—8.36 ppm (s, 1H)—7.59–6.69 (m, 10H)3.34 ppm (s, 2H)NiL1—8.31 ppm (s, 1H)—7.78–6.69 (m, 10H)—The1H-NMR spectrum of the ligand is presented in Figure 2. The phenolic proton of the ligand appears downfield as a singlet at 15.01 ppm while the signal at 3.35 ppm was attributed to the thiol (-SH) proton. The broadness of the signal was due to a strong hydrogen bonding between the imine nitrogen and the phenolic proton [34–36]. On the other hand, the azomethine proton (-HC=N-) appeared as a strong singlet at 9.78 ppm, confirming the formation of the Schiff base. All the aromatic protons were observed at 7.89–7.02 ppm [38].Figure 2
1H-NMR spectrum of the Schiff base ligand H2L1.A comparison of1H-NMR spectra of copper, cobalt, and nickel complexes with that of the free ligand displayed an upfield shift in the signal of azomethine proton which appeared at 8.09–8.36 ppm suggesting the coordination of azomethine nitrogen to the metal ions [37]. However, signals at 15.01 ppm and 3.35 ppm corresponding to a hydroxyl group and thiol group, respectively, in free ligand were found to be absent in both complexes, indicating the deprotonation of the hydroxyl group of naphthaldehyde moiety and thiol group of 2-aminothiophenol [38]. However, the peak in the range 10-11 ppm corresponds to the possible enamine tautomerization proton. Moreover, the signal observed at 3.34 ppm, corresponding to two protons in Co(II) complex spectrum, is assigned to one water molecule of crystallization [38]. The above conclusion is in agreement with the IR analysis which confirmed that the ligand is tridentate and coordinates to the metal using the azomethine nitrogen, phenolic oxygen, and thiophenolic sulfur atoms. The 1H-NMR data of the ligand and its complexes are presented in Table 3.


## 3.3. Mass Spectra

The ESI-MS of the Schiff base was performed to determine the molecular weight and fragmentation pattern. The molecular ion peak was observed atm/z 279.2 confirming formula weight (C17H13NOS) of H2L1. The peak at m/z 302.2 can be assigned to [M + Na]+ and the peak at m/z = 304.1 is assigned to [MH + Na]+ probable (13C-isotope in this case, due to its very low relative abundance) [39]. The mass spectrum of H2L1 is shown in Figure 3.Figure 3
Mass spectra of the Schiff base ligand H2L1.


## 3.4. Electronic Spectra

The UV-visible spectra of compounds are presented in Figure4, and their electronic data in DMSO are given in Table 4. The spectrum of H2L1 exhibited high-intensity bands at 362 nm (27624 cm−1), assigned to π ⟶ π∗ transition, and at 465 nm (21505 cm−1) assigned to n ⟶ π∗ transition [39, 40]. The band at 362 nm in the spectrum of the Schiff base experienced a blue shift in the spectra of the complexes. The spectrum of CuL1 complex shows an absorption band at 586 nm attributed to 2B1g ⟶ 2A1g transition of square planar geometry around Cu(II) while the band observed at 436 nm was attributed to charge transfer transition [40]. The spectrum of CoL1 complex shows an absorption band at 640 nm attributed to 4A2 ⟶ 4T1(P) transition which corresponds to tetrahedral geometry, while the band observed at 476 nm corresponds to charge transfer transition [41–44]. The spectrum of NiL1 showed absorption bands at 564 nm, attributed to 3T2(F) ⟶ 3T1(F) transition which corresponds to tetrahedral geometry, and the band observed at 453 nm is charge transfer transitions [39].Figure 4
Electronic spectrum of ligand H2L1 and its corresponding complexes.
(a)(b)(c)(d)Table 4
Electronic spectral data of the ligand and its corresponding complexes.
Compoundsλ (nm) andν (cm−1)ε (L·mol−1·cm−1)AssignmentsSuggested structuresH2L1362 nm (27624 cm−1)2901π ⟶ π∗—465 nm (21505 cm−1)740n ⟶ π∗CuL1332 nm (30120 cm−1)1610H2L1 chromophore (π ⟶ π∗)Square planar436 nm (22935 cm−1)1113Charge transfer (L ⟶ M)586 nm (17064 cm−1)402B1g ⟶ 2A1gCoL1339 nm (29498 cm−1)1137H2L1 chromophore (π ⟶ π∗)Tetrahedral476 nm (21598 cm−1)578Charge transfer (L ⟶ M)640 nm (15625 cm−1)1764A2 ⟶ 4T1 (P)NiL1329 nm (30395 cm−1)626H2L1 chromophore (π ⟶ π∗)Tetrahedral453 nm (22075 cm−1)772Charge transfer (L ⟶ M)564 nm (17730 cm−1)463T1(F) ⟶ 3T2 (F)


## 3.5. Thermal Analysis

The differential and thermogravimetric analyses of CuL1, CoL1, and NiL1 complexes were determined under a nitrogen gas inert atmosphere in the range 0 and 600°C. The thermograms and data obtained are shown in Figure 5 and Table 5, respectively. The thermal decomposition curves of CuL1 and NiL1 complexes showed single step decomposition corresponding to the loss of two molecules of ligand at 385°C (found: 72.11; calc: 72.56) and 428°C (found: 68.36; calc: 68.55), respectively, with a DTA of 400°C and 430°C, respectively. The thermal decomposition curve of the CoL1 complex showed two-step decomposition. The first step at 220°C gave a loss (found: 15.12; calc: 15.11) which corresponds to one molecule of aminothiophenol, while the second step occurred in the range of 400–556°C (found: 44.05; calc: 44.10) corresponding to the loss of one molecule of aminothiophenol and two molecules of 2-hydroxy-1-naphthaldehyde [39].Figure 5
TGA and DTA of CuL1, NiL1, and CoL1 complexes.
(a)(b)(c)Table 5
Thermoanalytical results for CuL1, CoL1, and NiL1 complexes.
CompoundsStepTGA temperature (°C)DTA temperature (°C)Weight loss (%)ProcessFragment lostFoundCalcCuL1Single385400 (endothermic)72.1172.56DecompositionC17H13NO (2 moles)Residue>60023.0323.34Cu2SCoL11220384 (endothermic)15.1215.11DecompositionC6H7NS (1 mole) + 3/4H2O2400–556457 (endothermic)44.0544.10DecompositionC11H8O2 (2 moles) +C6H7NS (1 mole)Residue>600—40,1241.52Co2O3; CoS; CONiL1Single428430 (endothermic)68.3668.55DecompositionC17H13NS (2 moles)Residue>60030.9130.96Ni3O2


## 3.6. Proposed Structures of Complexes

In agreement with microanalysis and all the spectroscopic results, the tridentate Schiff base coordinated to the metal using nitrogen atom of the azomethine group (-HC=N-), oxygen atom of hydroxynaphthaldehyde, and sulfur atom of the aminothiophenol. During the coordination, complexes probably dimerized to form neutral dimeric four-coordinate systems [33]. The dimerization is presumed to occur as a result of the bridging between the M(II) centers through the deprotonated thiophenolic sulfur atoms as described in Figure 6.Figure 6
Proposed structure of the M(II) complexes.


## 4. Cyclic Voltammetry Study

The electrochemical behavior of H2L1 and its Co(II), Ni(II), and Cu(II) complexes were investigated in DMF between −1.5 V and 1.5 V and a scan rate of 100 mV·s−1.


### 4.1. Cyclic Voltammetry Study of Ligand H2L1

The cyclic voltammograms of the ligand (Figure7) show three anodic waves at Epa1 = 0.587 V, Epa2 = 0.897 V, and Epa3 = 1.105 V. The first one, observed at 0.587 V, may be ascribed to the oxidation of the naphtone generated from the tautomeric equilibrium (Ketamine-Enamine) (see Scheme 2).Figure 7
Cyclic voltammogram of blank and ligand H2L1.Scheme 2
Scheme showing the possibility of the Schiff base to tautomerize [31].The last two waves at 0.897 and 1.105 V are attributed to the oxidation of thiophenolic and naphtholic groups, respectively [41–46]. For the return sweep, three cathodic waves were observed at 0.748, −0.141, and −1.075 V. The first wave observed at Epc1 = 0.748 V corresponds to the reduced species of the oxidized thiophenolic with naphtholic groups. The second wave is attributed to the naphtholic reduced form at Epc2 = −0.141 V while the third one observed at Epc3 = −1.075 V is attributed to the reduction of the azomethine group [41–46].


### 4.2. Cyclic Voltammetry Study of Cu(II) Complex

The cyclic voltammogram of CuL1 complex (Figure 8) exhibits four oxidation peaks. The first peak at Epa1 = 0.082 V corresponds to the oxidation of Cu(I) to Cu(II), and the last three peaks at Epa2 = 0.352 V, Epa3 = 0.943 V, and Epa4 = 1.174 V are attributed to the oxidation of azomethine group, thiophenolic group, and naphtholic group, respectively [44]. On the cathodic side, two peaks were observed at Epc1 = −0.346 V and Epc2 = −0.705 V which correspond to the reduction of Cu(II) to Cu(I) and the reduction of the azomethine group, respectively [46]. When this cyclic voltammogram is compared to that of the ligand, the shift in the reduction potential of the azomethine group is observed. This confirms the involvement of the azomethine group in coordination. However, the oxidation potential of the thiophenolic and naphtholic moieties seems to be more positive than the value observed for the ligand, which may be related to the relative stability of the coordination bonding between the oxygen and sulfur atom with copper (II) ion [45]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 3.Figure 8
Cyclic voltammograms of Schiff base H2L1 and Cu(II) complex.Scheme 3
Electron transfer process in Cu(II) complex.


### 4.3. Cyclic Voltammetry Study of Co(II) Complex

The cyclic voltammograms of CoL1 complex (Figure 9) exhibit one oxidation peak at Epa = −0.603 V corresponding to the oxidation of Co(II) to Co(III) and a reduction peak at Epc = −0.474 V corresponding to the reduction of Co(III) to Co(II) [44–46]. When this cyclic voltammogram is compared to that of the ligand, a disappearance of the azomethine group reduction peak in the complex is observed. However, the oxidation potential of the thiophenolic and naphtholic moieties seems to be more positive than the value observed for the ligand which may be related to the relative stability of the coordination bonding between the oxygen and sulfur atom with cobalt (II) ion [46]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 4.Figure 9
Cyclic voltammograms of Schiff base H2L1 and Co(II) complex.Scheme 4
Electron transfer process in Co(II) complex.


### 4.4. Cyclic Voltammetry Study of Ni(II) Complex

The cyclic voltammogram of NiL1 complex (Figure 10) exhibits one oxidation peak at Epa = 0.840 V corresponding to the oxidation of Ni(II) to Ni(III) and a reduction peak observed at Epc = 0.487 V corresponding to the reduction of Ni(III) to Ni(II) [46]. When this cyclic voltammogram is compared to that of the ligand, a disappearance of the azomethine group reduction peak in the complex is observed. However, the oxidation potential of the thiophenolic and naphtholic groups seems to be more positive than the value observed for the ligand [45]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 5.Figure 10
Cyclic voltammograms of Schiff base H2L1 and Ni(II) complex.Scheme 5
Electron transfer process in Ni(II) complex.In order to study the kinetics of the electrochemical process, the redox systems of the complexes were investigated separately in the reduced potentials range of −1 to 1.5 V using various scan rates (10 to 220 mV·s−1). The obtained voltammograms are given in Figure 11. The peak-to-peak separation ΔEp of 0.428 V and 0.592 V for Cu(II)/Cu(I) and Co(III)/Co(II), respectively, indicates the quasi-reversibility of the redox system studied. The ratio of cathodic to anodic peak height was close to 1. However, the peak current increases with an increase in the square root of the scan rates (Figure 12). This establishes the electrode process as diffusion controlled. The ratio of cathodic to anodic peak height of redox couple Ni(III)/Ni(II) is greater than 1, which indicates the irreversibility of the system. We can only note that the plot of the cathodic potential peak versus the logarithm of scan rate showed a nonlinear relationship (Figure 13) but the cathodic potential depends on the scan rate This indicates that the processes at the surface of the electrodes are slow [42–46]. The cyclic voltammograms of the ligand and its complexes were scanned 20 times at a scan rate of 10 mV·s−1 and superposed. Since all the peaks were observed in the same position in all the cycles, it can be concluded that there is no adsorption of the compound on the glassy carbon electrode surface [46].Figure 11
Scan rates of Cu(II) complex, Co(II) complex, and Ni(II) complex.
(a)(b)(c)Figure 12
Plot ofIp versus v1/2 for the (a) Cu(II)complex, (b) Co(II)complex, and (c) Ni(II)complex.
(a)(b)(c)Figure 13
Plot ofEpc versus LogV for the (a) Cu(II)complex, (b) Co(II)complex, and (c) Ni(II)complex.
(a)(b)(c)


## 4.1. Cyclic Voltammetry Study of Ligand H2L1

The cyclic voltammograms of the ligand (Figure7) show three anodic waves at Epa1 = 0.587 V, Epa2 = 0.897 V, and Epa3 = 1.105 V. The first one, observed at 0.587 V, may be ascribed to the oxidation of the naphtone generated from the tautomeric equilibrium (Ketamine-Enamine) (see Scheme 2).Figure 7
Cyclic voltammogram of blank and ligand H2L1.Scheme 2
Scheme showing the possibility of the Schiff base to tautomerize [31].The last two waves at 0.897 and 1.105 V are attributed to the oxidation of thiophenolic and naphtholic groups, respectively [41–46]. For the return sweep, three cathodic waves were observed at 0.748, −0.141, and −1.075 V. The first wave observed at Epc1 = 0.748 V corresponds to the reduced species of the oxidized thiophenolic with naphtholic groups. The second wave is attributed to the naphtholic reduced form at Epc2 = −0.141 V while the third one observed at Epc3 = −1.075 V is attributed to the reduction of the azomethine group [41–46].


## 4.2. Cyclic Voltammetry Study of Cu(II) Complex

The cyclic voltammogram of CuL1 complex (Figure 8) exhibits four oxidation peaks. The first peak at Epa1 = 0.082 V corresponds to the oxidation of Cu(I) to Cu(II), and the last three peaks at Epa2 = 0.352 V, Epa3 = 0.943 V, and Epa4 = 1.174 V are attributed to the oxidation of azomethine group, thiophenolic group, and naphtholic group, respectively [44]. On the cathodic side, two peaks were observed at Epc1 = −0.346 V and Epc2 = −0.705 V which correspond to the reduction of Cu(II) to Cu(I) and the reduction of the azomethine group, respectively [46]. When this cyclic voltammogram is compared to that of the ligand, the shift in the reduction potential of the azomethine group is observed. This confirms the involvement of the azomethine group in coordination. However, the oxidation potential of the thiophenolic and naphtholic moieties seems to be more positive than the value observed for the ligand, which may be related to the relative stability of the coordination bonding between the oxygen and sulfur atom with copper (II) ion [45]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 3.Figure 8
Cyclic voltammograms of Schiff base H2L1 and Cu(II) complex.Scheme 3
Electron transfer process in Cu(II) complex.


## 4.3. Cyclic Voltammetry Study of Co(II) Complex

The cyclic voltammograms of CoL1 complex (Figure 9) exhibit one oxidation peak at Epa = −0.603 V corresponding to the oxidation of Co(II) to Co(III) and a reduction peak at Epc = −0.474 V corresponding to the reduction of Co(III) to Co(II) [44–46]. When this cyclic voltammogram is compared to that of the ligand, a disappearance of the azomethine group reduction peak in the complex is observed. However, the oxidation potential of the thiophenolic and naphtholic moieties seems to be more positive than the value observed for the ligand which may be related to the relative stability of the coordination bonding between the oxygen and sulfur atom with cobalt (II) ion [46]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 4.Figure 9
Cyclic voltammograms of Schiff base H2L1 and Co(II) complex.Scheme 4
Electron transfer process in Co(II) complex.


## 4.4. Cyclic Voltammetry Study of Ni(II) Complex

The cyclic voltammogram of NiL1 complex (Figure 10) exhibits one oxidation peak at Epa = 0.840 V corresponding to the oxidation of Ni(II) to Ni(III) and a reduction peak observed at Epc = 0.487 V corresponding to the reduction of Ni(III) to Ni(II) [46]. When this cyclic voltammogram is compared to that of the ligand, a disappearance of the azomethine group reduction peak in the complex is observed. However, the oxidation potential of the thiophenolic and naphtholic groups seems to be more positive than the value observed for the ligand [45]. The controlled potential electrolysis carried out shows that the peak corresponds to a one-electron transfer process, as given in Scheme 5.Figure 10
Cyclic voltammograms of Schiff base H2L1 and Ni(II) complex.Scheme 5
Electron transfer process in Ni(II) complex.In order to study the kinetics of the electrochemical process, the redox systems of the complexes were investigated separately in the reduced potentials range of −1 to 1.5 V using various scan rates (10 to 220 mV·s−1). The obtained voltammograms are given in Figure 11. The peak-to-peak separation ΔEp of 0.428 V and 0.592 V for Cu(II)/Cu(I) and Co(III)/Co(II), respectively, indicates the quasi-reversibility of the redox system studied. The ratio of cathodic to anodic peak height was close to 1. However, the peak current increases with an increase in the square root of the scan rates (Figure 12). This establishes the electrode process as diffusion controlled. The ratio of cathodic to anodic peak height of redox couple Ni(III)/Ni(II) is greater than 1, which indicates the irreversibility of the system. We can only note that the plot of the cathodic potential peak versus the logarithm of scan rate showed a nonlinear relationship (Figure 13) but the cathodic potential depends on the scan rate This indicates that the processes at the surface of the electrodes are slow [42–46]. The cyclic voltammograms of the ligand and its complexes were scanned 20 times at a scan rate of 10 mV·s−1 and superposed. Since all the peaks were observed in the same position in all the cycles, it can be concluded that there is no adsorption of the compound on the glassy carbon electrode surface [46].Figure 11
Scan rates of Cu(II) complex, Co(II) complex, and Ni(II) complex.
(a)(b)(c)Figure 12
Plot ofIp versus v1/2 for the (a) Cu(II)complex, (b) Co(II)complex, and (c) Ni(II)complex.
(a)(b)(c)Figure 13
Plot ofEpc versus LogV for the (a) Cu(II)complex, (b) Co(II)complex, and (c) Ni(II)complex.
(a)(b)(c)


## 5. Biological Study

### 5.1. Antimicrobial Activity of the Ligand and Its Metal Complexes

The Schiff base and its metal complexes were tested against five bacterial strains (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, Enterococcus faecalis, and Proteus mirabilis) and five fungi (Candida albicans, Candida glabrata, Candida tropicalis, Candida krusei, and Candida parapsilosis) while Ciprofloxacin and Ketoconazole were used as the standard for bacterial and fungal studies, respectively.


#### 5.1.1. Determination of the Diameters of Zone of the Inhibition (DZI)

The diameters of the zone of inhibition of the ligand and its complexes were determined using the agar well diffusion method (Figure14) [29]. The diameters of zone of inhibition of the bacteria are summarized in Table 6 while Figure 15 is a histogram of the diameter of zone of inhibition.Figure 14
Antibacterial activity of Schiff base ligand(1) and Cu(II)complex(2), Co(II)complex(3), and Ni(II)complex(4) onStaphylococcus aureus.Table 6
Diameter of the zone of inhibition (DZI) of bacteria.
Diameter of zone of inhibition (DZI)BacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinE. coli——12—29S. aureus—1312—39P. aeruginosa———926E. faecalis————29P. mirabilis—10—1130Compounds are considered active when their diameters of the zone of inhibition are greater than 6 mm (DZI ˃ 6 mm) and (—) no activities are observed.Figure 15
Histogram representing the diameter of zone of inhibition (DZI) in mm of compounds.The results revealed that the synthesized compounds exhibited various levels of inhibition. It was found that the Schiff base H2L1 did not show any antibacterial activity against the tested bacterial and fungal strains except Candida parapsilosis on which activity is moderate while its metal(II) complexes showed moderate antibacterial activity compared to the ligand but very low compared to the standards used on Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, and Proteus mirabilis tested. This enhancement in the activity of the metal complexes can be explained on the basis of the chelation theory [47–50]. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups. Furthermore, the mode of action of the compound may involve the formation of a hydrogen bond through the azomethine group with the active center of cell constituents, resulting in interference with the normal cell process [50].


#### 5.1.2. Minimum Inhibitory Concentration (MIC)

The minimal inhibitory concentrations of the compounds were determined using the microdilution in a liquid environment in 96-well microtiter plates [31, 32, 51]. The microbial culture was placed in the presence of the compounds in decreasing order of concentration, in the wells of the microplates. After incubation, the lowest concentration of the antimicrobials in which there was no visible growth of the microorganism represents their minimum inhibitory concentration.


#### 5.1.3. Minimum Bactericidal Concentration (MBC)

The minimum bactericidal concentrations of the compounds were determined by sowing 10μL of the content of every well presenting no manifested change of color after the revelation at the p-iodonitrotetrazolium chloride (INT). After sowing in the sterile boxes of a Petri dish of 90 mm containing MHA, these boxes were incubated at 37°C for 24 hours. At the end of the incubation, the smallest concentration underneath of which no resumption of the bacterial growth was observed is the minimum bactericidal concentration. The minimum inhibitory concentration (MIC) and minimum bactericidal concentration (MBC) results are summarized in Table 7.Table 7
Minimum inhibition concentration (MIC) and minimum bactericidal concentration (MBC) (μg/mL).
bacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinMICMBCMICMBCMICMBCMICMBCMICMBCE. coli>256>25632256128>256128>256212S. aureus>256>256321286425664256112P. aeruginosa256>25664256128>2566425628E. faecalis256>256128>256>256>2566425644P. mirabilis256>25664256>256>2563212828


#### 5.1.4. Minimum Fungicidal Concentration (MFC)

The minimum fungicidal concentrations were obtained by planting out on liquid environment (every well of the 96-well microtiter plate containing 150μL of culture medium) 50 μL of the content of the wells that did not present any growth. The smallest concentrations that induced an absence of turbidity at the bottom of the wells after incubation were noted as the minimum fungicidal concentrations. The minimum inhibitory concentration (MIC) and minimum fungicidal concentration (MBC) results are summarized in Table 8.Table 8
Minimum inhibition concentration (MIC) and minimum fungicidal concentration (MFC) (μg/mL).
FungiCompoundsStandardH2L1CuL1CoL1NiL1KetoconazoleMICMBCMICMBCMICMBCMICMBCMICMBCC. albicans256>256256>256256>256256>2560.5064.00C. glabrata256>256>256>256>256>256>256>2560.258.00C. krusei256>256256>256256>256>256>2560.1250.50C. tropicalis256>256256>256>256>256>256>2568.008.00C. parapsilosis64256256>256>256>256>256>2562.0016.00In Table8, 256 μg/mL is the minimal value of the concentrations of antifungal and antibacterial in which there is no visible growth of the microorganism.The activities of compounds are considered significant when MIC and MFC or MBC < 10μg/mL, moderate when 10 ≤ MIC and MFC or MBC ≤ 256 μg/mL, and weak when MIC and MFC or MBC > 256 μg/mL [50, 51].It is observed from the biological data that the Schiff base presents weak antifungal activity, except againstC. parapsilosis where it shows moderate activity with MIC = 64 μg/mL, while the majority of complexes present weak antifungal activities with MIC > 256 μg/mL. The copper complex is weakly active against Staphylococcus aureus and Proteus mirabilis while the nickel complex is weakly active against Pseudomonas aeruginosa and Proteus mirabilis The cobalt complex shows weak activity against Escherichia coli and Staphylococcus aureus. However, these complexes are significantly more active than the ligand.


### 5.2. Antioxidant Activity of the Schiff Base and Its Metal Complexes

The synthesized compounds were screened for their free radical scavenging activities by the DPPH method using Trolox as a standard. Antioxidant activities of these compounds were investigated by measuring the radical scavenging effect on DPPH radicals. The results of the free radical scavenging activity of the compounds at different concentrations are shown in Figure16. It is evident from these results that the free radical scavenging activities of the compounds are concentration-dependent [49, 52].Figure 16
DPPH radical scavenging activity of the ligand and its metal complexes.The ligand exhibited good scavenging activity. All complexes showed moderate scavenging activity compared to the ligand. This observation could be due to free electron mobility in the complexes [51, 52] and deprotonation of ligand during chelation. The Cu(II) complex exhibited moderate scavenging activity compared to Co(II) and Ni(II) complexes.EC50 values for scavenging free radicals (Figure17) confirm the above affirmation that the ligand and its Cu(II) complex are more potent than the Ni(II) and Co(II) complexes.Figure 17
EC50 inhibitory concentration of the H2L1 ligand and its metal complexes for 50% of DPPH radical comparison was made against Trolox.


## 5.1. Antimicrobial Activity of the Ligand and Its Metal Complexes

The Schiff base and its metal complexes were tested against five bacterial strains (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, Enterococcus faecalis, and Proteus mirabilis) and five fungi (Candida albicans, Candida glabrata, Candida tropicalis, Candida krusei, and Candida parapsilosis) while Ciprofloxacin and Ketoconazole were used as the standard for bacterial and fungal studies, respectively.


### 5.1.1. Determination of the Diameters of Zone of the Inhibition (DZI)

The diameters of the zone of inhibition of the ligand and its complexes were determined using the agar well diffusion method (Figure14) [29]. The diameters of zone of inhibition of the bacteria are summarized in Table 6 while Figure 15 is a histogram of the diameter of zone of inhibition.Figure 14
Antibacterial activity of Schiff base ligand(1) and Cu(II)complex(2), Co(II)complex(3), and Ni(II)complex(4) onStaphylococcus aureus.Table 6
Diameter of the zone of inhibition (DZI) of bacteria.
Diameter of zone of inhibition (DZI)BacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinE. coli——12—29S. aureus—1312—39P. aeruginosa———926E. faecalis————29P. mirabilis—10—1130Compounds are considered active when their diameters of the zone of inhibition are greater than 6 mm (DZI ˃ 6 mm) and (—) no activities are observed.Figure 15
Histogram representing the diameter of zone of inhibition (DZI) in mm of compounds.The results revealed that the synthesized compounds exhibited various levels of inhibition. It was found that the Schiff base H2L1 did not show any antibacterial activity against the tested bacterial and fungal strains except Candida parapsilosis on which activity is moderate while its metal(II) complexes showed moderate antibacterial activity compared to the ligand but very low compared to the standards used on Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, and Proteus mirabilis tested. This enhancement in the activity of the metal complexes can be explained on the basis of the chelation theory [47–50]. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups. Furthermore, the mode of action of the compound may involve the formation of a hydrogen bond through the azomethine group with the active center of cell constituents, resulting in interference with the normal cell process [50].


### 5.1.2. Minimum Inhibitory Concentration (MIC)

The minimal inhibitory concentrations of the compounds were determined using the microdilution in a liquid environment in 96-well microtiter plates [31, 32, 51]. The microbial culture was placed in the presence of the compounds in decreasing order of concentration, in the wells of the microplates. After incubation, the lowest concentration of the antimicrobials in which there was no visible growth of the microorganism represents their minimum inhibitory concentration.


### 5.1.3. Minimum Bactericidal Concentration (MBC)

The minimum bactericidal concentrations of the compounds were determined by sowing 10μL of the content of every well presenting no manifested change of color after the revelation at the p-iodonitrotetrazolium chloride (INT). After sowing in the sterile boxes of a Petri dish of 90 mm containing MHA, these boxes were incubated at 37°C for 24 hours. At the end of the incubation, the smallest concentration underneath of which no resumption of the bacterial growth was observed is the minimum bactericidal concentration. The minimum inhibitory concentration (MIC) and minimum bactericidal concentration (MBC) results are summarized in Table 7.Table 7
Minimum inhibition concentration (MIC) and minimum bactericidal concentration (MBC) (μg/mL).
bacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinMICMBCMICMBCMICMBCMICMBCMICMBCE. coli>256>25632256128>256128>256212S. aureus>256>256321286425664256112P. aeruginosa256>25664256128>2566425628E. faecalis256>256128>256>256>2566425644P. mirabilis256>25664256>256>2563212828


### 5.1.4. Minimum Fungicidal Concentration (MFC)

The minimum fungicidal concentrations were obtained by planting out on liquid environment (every well of the 96-well microtiter plate containing 150μL of culture medium) 50 μL of the content of the wells that did not present any growth. The smallest concentrations that induced an absence of turbidity at the bottom of the wells after incubation were noted as the minimum fungicidal concentrations. The minimum inhibitory concentration (MIC) and minimum fungicidal concentration (MBC) results are summarized in Table 8.Table 8
Minimum inhibition concentration (MIC) and minimum fungicidal concentration (MFC) (μg/mL).
FungiCompoundsStandardH2L1CuL1CoL1NiL1KetoconazoleMICMBCMICMBCMICMBCMICMBCMICMBCC. albicans256>256256>256256>256256>2560.5064.00C. glabrata256>256>256>256>256>256>256>2560.258.00C. krusei256>256256>256256>256>256>2560.1250.50C. tropicalis256>256256>256>256>256>256>2568.008.00C. parapsilosis64256256>256>256>256>256>2562.0016.00In Table8, 256 μg/mL is the minimal value of the concentrations of antifungal and antibacterial in which there is no visible growth of the microorganism.The activities of compounds are considered significant when MIC and MFC or MBC < 10μg/mL, moderate when 10 ≤ MIC and MFC or MBC ≤ 256 μg/mL, and weak when MIC and MFC or MBC > 256 μg/mL [50, 51].It is observed from the biological data that the Schiff base presents weak antifungal activity, except againstC. parapsilosis where it shows moderate activity with MIC = 64 μg/mL, while the majority of complexes present weak antifungal activities with MIC > 256 μg/mL. The copper complex is weakly active against Staphylococcus aureus and Proteus mirabilis while the nickel complex is weakly active against Pseudomonas aeruginosa and Proteus mirabilis The cobalt complex shows weak activity against Escherichia coli and Staphylococcus aureus. However, these complexes are significantly more active than the ligand.


## 5.1.1. Determination of the Diameters of Zone of the Inhibition (DZI)

The diameters of the zone of inhibition of the ligand and its complexes were determined using the agar well diffusion method (Figure14) [29]. The diameters of zone of inhibition of the bacteria are summarized in Table 6 while Figure 15 is a histogram of the diameter of zone of inhibition.Figure 14
Antibacterial activity of Schiff base ligand(1) and Cu(II)complex(2), Co(II)complex(3), and Ni(II)complex(4) onStaphylococcus aureus.Table 6
Diameter of the zone of inhibition (DZI) of bacteria.
Diameter of zone of inhibition (DZI)BacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinE. coli——12—29S. aureus—1312—39P. aeruginosa———926E. faecalis————29P. mirabilis—10—1130Compounds are considered active when their diameters of the zone of inhibition are greater than 6 mm (DZI ˃ 6 mm) and (—) no activities are observed.Figure 15
Histogram representing the diameter of zone of inhibition (DZI) in mm of compounds.The results revealed that the synthesized compounds exhibited various levels of inhibition. It was found that the Schiff base H2L1 did not show any antibacterial activity against the tested bacterial and fungal strains except Candida parapsilosis on which activity is moderate while its metal(II) complexes showed moderate antibacterial activity compared to the ligand but very low compared to the standards used on Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa, and Proteus mirabilis tested. This enhancement in the activity of the metal complexes can be explained on the basis of the chelation theory [47–50]. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups. Furthermore, the mode of action of the compound may involve the formation of a hydrogen bond through the azomethine group with the active center of cell constituents, resulting in interference with the normal cell process [50].


## 5.1.2. Minimum Inhibitory Concentration (MIC)

The minimal inhibitory concentrations of the compounds were determined using the microdilution in a liquid environment in 96-well microtiter plates [31, 32, 51]. The microbial culture was placed in the presence of the compounds in decreasing order of concentration, in the wells of the microplates. After incubation, the lowest concentration of the antimicrobials in which there was no visible growth of the microorganism represents their minimum inhibitory concentration.


## 5.1.3. Minimum Bactericidal Concentration (MBC)

The minimum bactericidal concentrations of the compounds were determined by sowing 10μL of the content of every well presenting no manifested change of color after the revelation at the p-iodonitrotetrazolium chloride (INT). After sowing in the sterile boxes of a Petri dish of 90 mm containing MHA, these boxes were incubated at 37°C for 24 hours. At the end of the incubation, the smallest concentration underneath of which no resumption of the bacterial growth was observed is the minimum bactericidal concentration. The minimum inhibitory concentration (MIC) and minimum bactericidal concentration (MBC) results are summarized in Table 7.Table 7
Minimum inhibition concentration (MIC) and minimum bactericidal concentration (MBC) (μg/mL).
bacteriaCompoundsStandardH2L1CuL1CoL1NiL1CiprofloxacinMICMBCMICMBCMICMBCMICMBCMICMBCE. coli>256>25632256128>256128>256212S. aureus>256>256321286425664256112P. aeruginosa256>25664256128>2566425628E. faecalis256>256128>256>256>2566425644P. mirabilis256>25664256>256>2563212828


## 5.1.4. Minimum Fungicidal Concentration (MFC)

The minimum fungicidal concentrations were obtained by planting out on liquid environment (every well of the 96-well microtiter plate containing 150μL of culture medium) 50 μL of the content of the wells that did not present any growth. The smallest concentrations that induced an absence of turbidity at the bottom of the wells after incubation were noted as the minimum fungicidal concentrations. The minimum inhibitory concentration (MIC) and minimum fungicidal concentration (MBC) results are summarized in Table 8.Table 8
Minimum inhibition concentration (MIC) and minimum fungicidal concentration (MFC) (μg/mL).
FungiCompoundsStandardH2L1CuL1CoL1NiL1KetoconazoleMICMBCMICMBCMICMBCMICMBCMICMBCC. albicans256>256256>256256>256256>2560.5064.00C. glabrata256>256>256>256>256>256>256>2560.258.00C. krusei256>256256>256256>256>256>2560.1250.50C. tropicalis256>256256>256>256>256>256>2568.008.00C. parapsilosis64256256>256>256>256>256>2562.0016.00In Table8, 256 μg/mL is the minimal value of the concentrations of antifungal and antibacterial in which there is no visible growth of the microorganism.The activities of compounds are considered significant when MIC and MFC or MBC < 10μg/mL, moderate when 10 ≤ MIC and MFC or MBC ≤ 256 μg/mL, and weak when MIC and MFC or MBC > 256 μg/mL [50, 51].It is observed from the biological data that the Schiff base presents weak antifungal activity, except againstC. parapsilosis where it shows moderate activity with MIC = 64 μg/mL, while the majority of complexes present weak antifungal activities with MIC > 256 μg/mL. The copper complex is weakly active against Staphylococcus aureus and Proteus mirabilis while the nickel complex is weakly active against Pseudomonas aeruginosa and Proteus mirabilis The cobalt complex shows weak activity against Escherichia coli and Staphylococcus aureus. However, these complexes are significantly more active than the ligand.


## 5.2. Antioxidant Activity of the Schiff Base and Its Metal Complexes

The synthesized compounds were screened for their free radical scavenging activities by the DPPH method using Trolox as a standard. Antioxidant activities of these compounds were investigated by measuring the radical scavenging effect on DPPH radicals. The results of the free radical scavenging activity of the compounds at different concentrations are shown in Figure16. It is evident from these results that the free radical scavenging activities of the compounds are concentration-dependent [49, 52].Figure 16
DPPH radical scavenging activity of the ligand and its metal complexes.The ligand exhibited good scavenging activity. All complexes showed moderate scavenging activity compared to the ligand. This observation could be due to free electron mobility in the complexes [51, 52] and deprotonation of ligand during chelation. The Cu(II) complex exhibited moderate scavenging activity compared to Co(II) and Ni(II) complexes.EC50 values for scavenging free radicals (Figure17) confirm the above affirmation that the ligand and its Cu(II) complex are more potent than the Ni(II) and Co(II) complexes.Figure 17
EC50 inhibitory concentration of the H2L1 ligand and its metal complexes for 50% of DPPH radical comparison was made against Trolox.


## 6. Conclusion

We have prepared and characterized a novel tridentate Schiff base ligand (H2L1) which coordinates easily to Co(II), Cu(II), and Ni(II) ions to form tetrahedral CoL1, NiL1, and square planar CuL1 complexes. Cyclic voltammetry studies on the ligand and its metal complexes reveal that the redox systems Co(III)/Co(II) and Cu(II)/Cu(I) displayed quasi-reversible processes while Ni(III)/Ni(II) displayed an irreversible process. Antibacterial studies show that the complexes were more active than the free ligand on some bacterial strains while antifungal studies reveal that these compounds showed no antifungal activity. Results of the antioxidant study showed that the ligand and its Cu(II) complex are more efficient to prevent the formation of the DPPH radical than Co(II) and Ni(II) complexes.



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*Source: 5238501-2020-11-09.xml*