The chemicals, 1-phenylsemicarbazide (Riedel-de Haën), 2-methoxybenzaldyde (Aldrich), diphenyl chlorophosphate (Riedel-de Haën), Triethyl amine (Riedel-de Haën), bipyridine (Aldrich), ferric chloride (Aldrich), ascorbic acid (BDH), sodium acetate trihydrate (BDH), glacial acetic acid (BDH) and metals chloride hydrate (CuCl₂.2H₂O, NiCl₂.6H₂O, and CoCl₂.6H₂O) were purchased from BDH. The solvents used were of spectroscopic grade.

Elemental analysis for C, H, and N was performed with a Vario EL Fab. CHN Nr, at Central Laboratory, Faculty of Science, Cairo University, Giza, Egypt. Conductivity measurements were done with a Jenway conductivity meter model 4510 in dimethylformamide (DMF). The magnetic moments were determined by Gouy’s method using a magnetic susceptibility balance from Johnson Metthey and Sherwood model. Ultraviolet visible (UV-vis) spectra were measured with (Specord200, Analytilk Jena, Germany) in the range of 200–900 nm, at Sana’a University. Fourier-transform infrared (FT-IR) spectra were recorded in transmission mode by an (FT/IR-140, Jasco, Japan) spectrophotometer in the wavenumber range of 4000–400 cm^–1 (KBr was used as a matrix material for pellets). ¹H and ¹³C NMR spectra for ligands were performed in d6-DMSO solvent using tetramethylsilane as an internal standard, at Cairo University, Giza, Egypt. All melting points reported for the compounds are measured in glass capillary tubes in Celsius degrees. Chloride was determined gravimetrically by silver nitrate (Vogel, 1961). The amount of coordinated and uncoordinated water molecules was determined gravimetrically using the weight loss method (Vogel, 1961). The X-ray diffraction (XRD) patterns were obtained using XD-2 (Shimadzu ED-720) powder X-ray diffractometer at a voltage of 35 kV and a current of 20 mA using CuK(α) radiation in the range of 5° ˂ 2θ ˂ 70° at 1° min^–1 scanning rate and a wavelength 1.54056 Å at Yemen Geological Survey and Mineral Resources Board.

The molar ratio method was described by the Molar ratio method (Yoe and Jones, 1944). The concentrations of metal ions (Cu²⁺, Ni²⁺, and Co²⁺) were kept constant at (1×10^–3 mol L^–1) in methanol and the concentration of ligands (L₁ and L₂) was regularly varied in methanol. The absorbance of the prepared solutions was measured at the constant wavelength (λ_max). The absorbance values were plotted versus the molar ratio [ligand] / [metal ion]. The intersections of the obtained straight lines indicate the molar ratio of the stable complexes.

The percentage of crystallinity, XC (%) was calculated based on the integrated peak areas of the principal peaks (Shah et al., 2006). The crystallinity of the complexes is calculated relative to the crystallinity of the ligands as a ratio (Eq. 1):

(1)

whereA_complexandA_ligandare the areas under the principal peaks of the complex and ligand sample, respectively.

X-ray diffraction was also used to determine the average particle size (D) which was estimated by the Scherrer equation (Akhtar et al., 2015; Patterson, 1939) (Eq. 2):

(2)

where K is Scherrer constant and equals 0.94, λ is the X-ray wavelength of Cu-Kα radiations (1.5405 Å), β is full width at half maximum (FWHM) and θ is Bragg diffraction angle in degrees.

The total antioxidant activity of the compounds has been studied using ferric-bipyridine reducing capacity (FBRC) (Al-Azab et al., 2023). It was taken (1 mL, 10^–2 mol L^–1 FeCl₃.6H₂O) to 10 mL volumetric flask different volumes of the standard antioxidant ascorbic acid (0.1 g L^–1) were added (0.01, 0.02, 0.04, 0.06, 0.08, 0.1, 0.2 mL). This was followed by 2.0 mL 0.3 mol L^–1 acetate buffer (pH 4) and 1.0 mL bipyridine (6.4 × 10–3 mol L^–1). The volume was completed to the mark with deionized water. After 10 min of incubation at room temperature, the absorbance was recorded against a blank at 535 nm. The absorbance values were plotted against the concentration of the various antioxidant solutions (Fig. 3). Similarly, 0.02 mL of (0.1 mg mL^–1 methanol) of each tested compound was reacted with 1 mL, 10^–2 mol L^–1 FeCl₃.6H₂O solution, (6.4 × 10–3 mol L^–1) bipyridine and 2.0 mL 0.3 mol L^–1 acetate buffer (pH 4). This mixture was diluted to 10 mL with deionized water for the antioxidant assay. The fixed reaction time and fixed state measurement method were used to find out the antioxidant activity of these compounds in methanol as a solvent. All spectrophotometric measurements were measured using UV-vis spectrophotometer (Specord200, Analytikjena, Germany) at Sana’a University.

Figure 3.
Calibration curve of ascorbic acid with Fe-Bp complex.

The synthesis compounds were exanimated for their antimicrobial activity in the Laboratory of Microbiology at Sana’a University. They Germinated four types of bacteria (Staphylococcus aureus, Bacillus subtilis,Escherichia coli, and Pseudomonas aeruginosa) and one fungus (Candida albicans) on nutrient agar and Sabouraud agar solid media respectively. Using the filter paper disc method (Ericsson et al., 1960), it was prepared one concentration from each compound (1000 µg mL^–1) in DMSO. Then added (100 µL) from each prepared compound on a filter paper (Whatman No. l filter paper, 5 mm diameter) which contained the bacteria with agar solid media, all the Petri dishes were put in an incubation period at 37 °C for 24 h. Gentamicin 120 μg mL^–1 was used as a reference substance for bacteria and Mycostatin 30 µg mL^–1 for fungi. The results were registered by calculating the diameter of the inhibition zone (mm).

Important IR bands of the ligands and their complexes are given in Table 2 and Fig. 5a. The presence of IR bands at 1,655 and 1,602 cm^–1 of ligand L1 assigned to υ(C=O) and υ(C=N) respectively (Hossain et al., 2019). The strong bands at 3,334 and 3,397 cm^–1 are assignable and attributed to υ(NH) of this ligand (Nakamoto, 1998). The bands showed at 3,057, 2,958, and 1,267 cm^–1 belong to aromatic υ(C-H), aliphatic υ(C-H), and υ(-C-N) stretching mode of vibrations, respectively (Jassem et al., 2013). The band at 2,853 cm^–1 in its spectrum can be appointed to υ(-O-CH3) (Mohapatra et al., 2011).

The IR spectra of the complexes (Fig. 5a) are characterized by the appearance of broad bands at 3,424, 3,346, and 3,324 cm^–1 due to the water molecules in L1(Cu), L1(Ni), and L1(Co) complexes, respectively (González-García et al., 2016). In the spectra of all the complexes, no change in the frequency of (C=O) but the 1,602 cm^–1 band was shifted to a lower frequency indicating the involvement of the N atom of the azomethine group in coordination (Table 2). The bands of υ(NH, NH) disappeared in complexes due to the broad bands of water. New vibrations at 400–600 cm^–1 that are not present in the free ligand are attributed to the existence of υ(M-N) and υ(M-NH) (Yassin et al., 2020). Based on the observations, it has been suggested that the ligand L1 acts as bidentate.

Although the presence evidence of υ(M-Cl) could not be brought in the IR data due to instrumental limitation, the insolubility of L1 complexes in water and their non-electrolytic nature evinced that 2Cl^- coordinated with metal in L1 complexes.

In the ligand L2 spectrum (Fig. 5b), the bands observed at 1,654 and 1,591 cm^–1 corresponds to υ(C=O) and υ(C=N), respectively (Hossain et al., 2019). The presence of υ(P=O) and υ(P-O-C) is indicated by the appearance of bands at 1,188 and 1,072 cm^–1, respectively (El-khazandar, 1997).

The coordinated water molecules with metal ions in the complexes were proved by the bands 3,347, 3,407, and 3,312 (Aderoju and Sherifah, 2015). The IR spectra of the metal complexes of ligand L2 (Fig. 5b) proposed that this ligand tridentate with the phosphoryl-oxygen, amine-nitrogen, and azomethine-nitrogen due to the shift in the position of υ(P=O), υ(-NH) and υ(C=N) (Table 2).

Table 2.
The main IR absorption bands of the ligands and their complexes.

Figure 5.
Infrared spectra. (a) L1 and its complexes; (b) L2 and its complexes.

The absorption spectra of the ligands L1 and L2 and their metal complexes were recorded in DMSO in the range of 200–900 nm (Table 3).

The ligands L1 and L2 exhibited two absorption bands (Figs. 6a and b) at 37,037.03, 35,714.28 and 33,333.33, 32,258.06 cm^–1 that have been allocated for the π-π* and n-π* transitions, respectively (Mohammed et al., 2019). In the metal complexes, the transition of n-π* has been seen but they are shifted to another range. The shift in the transitions has been accounted for by the complexation between ligands with the metal ions.

The Cu complexes of L1 and L2 exhibited single broad absorption peaks at 13,513.51 and 13,333.33 cm^–1 (Figs. 6a and b), respectively, attributable to the 2Eg→2T2g transition, indicating a distorted octahedral structure (Al-Maydama et al., 2008; Gup and Kirkan, 2005).

The electronic spectra of the L1(Ni) and L2(Ni) complexes (Figs. 5a and b) showed absorption bands at 18,181.81, 15,151.51 and 17,543.85, 15,151.51 cm^–1, respectively, attributable to the 3A2g → 3T1g(P) and 3A2g → 3T1g(F) transitions, which is compatible with this complex having the octahedral structure (Gup and Kirkan, 2005).

The electronic spectra of the L1 and L2 complexes with Co (Figs. 6a and b) exhibited characteristic bands at 20,833.33, 15,873.01, 13,888.88 and 21,739.13, 20,000, 17,543.85 cm^–1 assigned to 4T1g→ 4T1g(P), 4T1g→ 4A2g and 4T1g→4T2g(F) transitions. These bands are associated with the octahedral structures (Gup and Kirkan, 2005).

This information with the effective magnetic moment (µeff) data (Table 3) of all complexes helped to support the suggested octahedral geometry (Al-Hakimi et al., 2011; Fouda et al., 2008).

Table 3.
Electronic spectral and magnetic moment data of the ligands and their complexes.

Figure 6.
UV-Vis spectra. (a) L1 ligand and its complexes; (b) L2 ligand and its complexes.

Investigation of the molecular structure of the complexes formed between the metal ions of (Cu²⁺, Ni²⁺, and Co²⁺) with ligands in L1 and L2 in methanol using molar ratio (Figs. 7 and 8) revealed the formation of (1:2) (M:L) complexes for L1 and (1:1) (M:L) complexes for L2 under investigation. The results are depicted in Table 4.

Figure 7.
Molar ratio plot for the complexes between metal ions of (a) Cu, (b) Ni and (c) Co with L1 in methanol.

Figure 8.
Molar ratio plot for the complexes between metal ions of (a) Cu, (b) Ni and (c) Co with L2 in methanol.

Table 4.
Molar ratios for determination of stoichiometry of the L1 and L2complexes in methanol.

The XRD patterns of the L1 and its complexes are shown in Fig. S5a–d (see Supplementary Material). The decrease of the XRD peak intensities of L1 complexes can be attributed to the reduction in crystallinities due to the complexation. This is in good agreement with the lowering in relative crystallinity calculated for these complexes (the crystallinity calculation in Table 5 is based on the integrated area of principal peaks of the complex to the L1 ligand obtaining a relative crystallinity) (Shah et al., 2006). The particle size calculated for L1 and its complexes (2.074–5.552 nm) in the range of nanoparticle size—the nanoparticle size reported in the literature ranges between 1–100 nm (Boverhof et al., 2015).

The particle size of L2 ligand was found within nanometric range (10.555 nm), nevertheless its complexes appeared as an amorphous character (Fig. S5e–h, see Supplementary Material). From this change, it is expected to improved properties of L2 complexes as compared with L2 ligand.

Table 5.
Data of the principal values of intensity of the ligands and L1 complexes from XRD spectra.

The ligands and L2 complexes were investigated using FBRC technique to determine the antioxidant activity (Al-Azab et al., 2023).

Fe(II)-bipyridine complex produced by the reaction of Fe(III) with antioxidant followed by bipyridine exhibited maximum absorption at 535 nm (Fig. 9). Total antioxidant activity is based on the redox reaction between compounds and Fe(III) at room temperature. The initial antioxidant concentration is indicated by the concentration of the oxidizing Fe(III).

Figure 9.
Stoichiometry outline of FBRC method.

The results in Table 6 showed that these compounds were found to possess high potent antioxidant activity due presence of a combination of donor sites such as amide oxygen and imine nitrogen (Kostova and Saso, 2013).

Table 6.
FBRC values express antioxidant activity as mg L–1 for prepared compounds to 1 mg L–1 of ascorbic acid

The synthetic ligands and their metal complexes havbeen examined for their antibacterial and antifungal activities using the disk diffusion method (Ericsson et al., 1960) against four types of Bacteria (Staphylococcus aureus,Bacillus subtilis, Escherichia coli, and Pseudomonas aeruginosa) and one fungus (Candida albicans) on nutrient agar and Sabouraud dextrose agar (SDA) solid media, respectively. The inhibition zone diameter of the compounds is shown in Table 7 and Fig. 10a–d. From these results, it was noted that:

1- No inhibition zone was observed for ligands and their metal complexes against the fungus (Candida albicans);

2- The metal complexes are more potent bactericides than the ligands;

3- The order of antibacterial activity for the compounds with Bacillus subtilis was L1(Cu) = L2(Ni) > L1(Ni) > L2(Cu) > L2 > L1 > L1(Co) = L2(Co), with Staphylococcus aureus was L2(Ni) > L1(Cu) > L1(Co) > L2(Co) > L2 > L1 > L1(Ni) = L2(Cu), with Escherichia Coli was L1(Co) > L1(Cu) > L2(Cu) > L2(Co) > L2 > L1 > L1(Ni) = L2(Ni) and with Pseudomonas aeruginosa was L1(Co) = L2(Co) > L1(Cu) > L2(Ni) > L2 > L1 > L1(Ni) = L2(Cu).

In general, the complexes showed good antibacterial compared to the free ligands. Based on the chelation theory (Franklin and Snow, 1989), the increased inhibition activity of the metal complexes may be explained on the basis that their structures mainly possess C=N bonds. Besides, the coordination decreases the polarity of the metal ion due to the partial sharing of its positive charge with donor groups and possible π-electron delocalization inside the chelate ring-shaped during coordination that makes the complexes more lipophilic. This increased lipophilicity enhances the penetration of the metal complexes into lipid membranes, blocks the metal binding sites in the enzymes, and limits the further development of the organisms.

Table 7.
Biological activities of the ligands and their metal complexes against bacteria and fungus (zone of inhibition in mm).

Figure 10.
Antibacterial activity of the ligands and their metal complexes against gram-positive bacterium (a) Bacillus subtilis; (b) Staphylococcus aureus; (c) Pseudomonas aeruginosa; and gram-negative bacterium (d) Escherichia coli.

y.jamil@su.edu.ye