Aminothiazole-Linked Metal Chelates: Synthesis, Density Functional Theory, and Antimicrobial Studies with Antioxidant Correlations.

During the current study, the new aminothiazole Schiff base ligands (S1 ) and (S2 ) were designed by reacting 1,3-thiazol-2-amine and 6-ethoxy-1,3-benzothiazole-2-amine separately with 3-methoxy-2-hydroxybenzaldehyde in good yields (68-73%). The ligands were characterized through various analytical, physical, and spectroscopic (FT-IR, UV-Vis, 1H and 13C NMR, and MS) methods. The ligands were exploited in lieu of chelation with bivalent metal (cobalt, nickel, copper, and zinc) chlorides in a 1:2 (M:L) ratio. The spectral (UV-Vis, FT-IR, and MS), as well as magnetic, results suggested their octahedral geometry. The theoretically optimized geometrical structures were examined using the M06/6-311G+(d,p) function of density function theory. Their bioactive nature was designated by global reactivity parameters containing a high hardness (η) value of 1.34 eV and a lower softness (σ) value of 0.37 eV. Different microbial species were verified for their potency (in vitro), revealing a strong action. The Gram-positive Micrococcus luteus and Gram-negative Escherichia coli gave the highest activities of 20 and 21 mm for compounds (8) and (7), respectively. The antifungal activity against the Aspergillus niger and Aspergillus terreus species gave the highest activities of 20 and 18 mm for compounds (7) and (6), respectively. The antioxidant activity, evaluated as DPPH and ferric reducing power, gave the highest inhibition (%) as 72.0 ± 0.11% (IC50 = 144 ± 0.11 μL) and 66.3% (IC50 = 132 ± 0.11 μL) for compounds (3) and (8), respectively. All metal complexes were found to be more biocompatible than free ligands due to their chelation phenomenon. The energies of LUMOs had a link with their activities.

Highlights

A facile synthesis of aminothiazole-containing Schiff bases and their metal chelates.

Combined experimental and theoretical exploration to decide their medicinal role.

Geometry optimization studies at the M06/6-311G+(d,p) level of density function theory.

The relative order of medicinal studies was as follows: reference drugs > metal chelates > ligands.

Introduction

The discovery of cisplatin has led the scientist to a continuous pursuit for developing novel and efficient metallo-drugs having lower side effects.[1] After its invention in the 1960s, metal-based complexes indeed earned a great deal of interest from different chemotherapeutic scientists.[2] The assessment of novel materials as antimicrobial drugs has been of significant importance due to the presence, and subsequently the development, of multidrug resistance to the widespread pathogens.[3] The diverse range of coordination abilities for metals is also considered to be important biologically,[4] industrially,[5] and medicinally.[6] Presently, several coordination complexes possessing antimicrobial activities are reported.[7] The literature revealed an enhancement of antibacterial activity for metal complexes than their respective ligands.[8] The prosthetic role of metals for metal complexes may be determined during their improved medicinal effects.[9] The research and development regarding the development of coordination chemistry has established a definite relationship for metallo-compounds and biological roles for these compounds.[10] The traditional organic drugs have no answer for the ever-increasing pathogenic resistance as well as their specificity in their mechanism of reaction.[11]

The aminothiazole derivatives are, by far, the most significant heterocycles,[12] which are widespread and essential to the diversity of natural products[13] and medicinal agents.[14] They have a range of therapeutic activities with a high degree of structural versatility that has been proved valuable in the quest for new therapeutic agents.[15] The extensive variety of the pharmacological actions of specific derivatives has shown that such compounds are also of undoubted importance.[16] Metal-based compounds are considered potential candidates for the replacement of traditionally employed organic drugs as several metal complexes of benzothiazoles have established their unique and potent behavior in medicine such as antimicrobial,[17] antileukemic,[18] anti-inflammatory,[19] and antidiabetic agents.[20] So, these things demand that there should be versatile changes in the structure and activity of such medicinally important compounds to cope with these challenges.[21] To overcome these drawbacks, the metal-based compounds are being considered a good replacement. Their higher yields, specificity in their reactions, and the increase in biological activities upon chelation have provoked the scientists to continue such work in the future.[22]

The pharmacological and structural properties of newly synthesized metal complexes resulting from two ligands, 2-methoxy-6-{[(1,3-thiazol-2-yl)imino]methyl}phenol and 2-{[(6-ethoxy-1,3-benzothiazol-2-yl)imino]methyl}-6-methoxyphenol, were investigated in this report using a combination of experimental and simulated insights (Scheme ). The biological activities of all these metal complexes of Schiff bases such as antibacterial, antifungal, and anticancer activities have also been taken into account.

Scheme 1

Synthesis of Thiazole-Based Ligands (S) and (S) and Their Metal Complexes (1)–(8)

Results and Discussion

An aldehyde compound, 3-methoxy-2-hydroxybenzaldehyde, was independently reacted with two aminothiazole moieties, 2-methoxy-6-{[(1,3-thiazol-2-yl)imino]methyl}phenol and 6-ethoxy-1,3-benzothiazol-2-amine, to produce two ligands, i.e., 2-methoxy-6-{[(1,3-thiazol-2-yl)amino]methyl}phenol (S) and 2-{[(6-ethoxy-1,3-benzothiazol-2-yl)imino]methyl}-6-methoxyphenol (S). The solid condensation products of the above chemical components were air-stable and dissolvable in acetonitrile, dioxane, ethanol, methanol, N,N-dimethylformamide (DMF), and dimethyl sulfoxide (DMSO). The ligands and subsequent metal complexes were also generated in the neutral pH range. The ligands were established to synthesize their transition complexes (1)–(8) by reacting with metal salts. Their microanalytical, physical, and spectral findings categorized the ligands and metal complexes, whereas the geometry of metal complexes (1)–(8) was determined by their elemental, LC–MS, FT-IR, and UV–Vis data. Their geometry was demonstrated to contain an octahedral arrangement. The physical attributes and microanalytical results for ligands and their metal complexes (1)–(8) are enlisted in Table S1. The color shift between ligands and metal complexes supported a ligand’s interaction with a metal. The metal complexes had a wider range of decomposition points when compared to their respective ligands. These transition metals remain substantially higher at melting points compared to the main group elements. This is the result of a strong connection of metals to transition metals due to the electrons being relocated, and the electrons are made available from both d- and s-orbitals.[23] In conjunction with various spectral techniques, microanalytical observations were used to produce their proposed molecular formulas. Their low conductance values have indicated their non-electrostatic existence. The magnetic susceptibility ranges verified the probable geometrical agreement of ligands across ions (Table S2).

FT-IR Spectra

The Fourier transform infrared spectra of ligands and metal complexes clearly indicated their binding mode (Figures S1–S8). The ligands (S) and (S) showed the taking out of a band at 3320 cm–1, which transitioned to a newly formed azomethine v(HC=N) functional group around 1632–1634 cm–1, which indicated that the aldo carbonyl group was condensed with the amine unit of the aminothiazole.[24] Due to stretching v(C–S) and v(C=N) vibrations of the thiazole ring and the aromatic carbonyl v(C=O), all ligands showed significant bands at 845–862, 1612–1615, 1725–1749, and 3075–3149 cm–1 correspondingly.[25] The ligands had the bands for v(C–H) and v(C=C) around 2928–2974 and 1540–1568 cm–1, respectively. Additionally, owing to the bending v(C–O) vibrations, the ligands had bands at the 1388–1392 cm–1 range.[26] With the following data displayed for metal complexes, the band that was visible on 1633–1653 cm–1 due to the azomethine link was changed to lesser frequencies of 11–13 cm–1 at 1611–1614 cm–1, indicative of the role of the azomethine group nitrogen in the complexation.[27] The aminothiazole moiety v(C=N) band was shifted from 1612–1615 cm–1 to 15–25 cm–1, indicating the existence of a metal–nitrogen (M–N) bond in metal complexes.[28] The aldehydic bands of v(C=O) were moved to about 15–25 cm–1 from 1725–1749 cm–1 to 1703–1726 cm–1, demonstrating the coordination of the aminothiazole nitrogen to metal ions. The appearance of a band at 527–532 cm–1 revealed the coordination of thiazole ring nitrogen to metal ions (M–O). The bands in the spectra at 2968–2972, 844–860, and 1542–1564 cm–1, labeled with aromatic v(C–H), v(C–S), v(C–O), and v(C=C) units, were intact, which indicated that they did not play a role in the complexation proccess.[29]

1H NMR Spectra

The proton nuclear magnetic resonance (1H NMR) spectral data further confirmed the chemical composition and structure of the synthesized molecules (Figures S9 and S10). All of the ligand protons presented their signals in the predicted places. The singlet and doublet peaks were observed in the ranges of 11.05–11.15 and 6.91–6.93 ppm correspondingly due to C1-NH and C3-H. The doublet and triplet peaks for methylene and methoxy protons were present at 3.97–3.96 and 1.33 ppm, respectively.[30] In the ranges of 7.26–7.59 and 7.22–7.55 ppm, the aromatic protons C13-H and C15-H of the benzothiazole group (ligand S) were detected as a doublet and singlet, respectively, attributed to the existence of an additional electron-withdrawing methoxy (OCH3) group. Their protonic, C13-H and C15-H, values were deshielded to be recorded as a single and doublet at 8.68 and 8.25 ppm accordingly. The C4-H and C6-H protons appeared as triplets and doublets at 7.06 and 6.77 ppm, respectively. C5-H exhibited a triplet peak in the ligand at 7.58–7.61 ppm.[31] The C12-H proton was recognized as a doublet for all ligands in the 7.49–7.59 ppm range. The aromatic protons C13-H and C15-H of the benzothiazole group were found as a doublet and singlet in ligands at 7.22–7.55 ppm, respectively.

13C NMR Spectra

The carbon nuclear magnetic resonance (13C NMR) spectral studies in the DMSO solvent delivered a strong sign to establish the proposed designs of ligands (Figures S11 and S12). The 13C NMR spectra revealed distinguishing peaks inside the anticipated range. For the ligands, the carbon C8 peaks of azomethine (C=N) appeared at 183.80–184.86 ppm, which intensely reinforced the creation of ligands by a condensation response. The C1 carbon was detected downfield at 159.63–159.83 ppm due to the inductive impact of nitrogen and the electronegative influence of oxygen, showing the presence of C=O carbon.[32] Carbon atom C9 existing amid the nitrogen and sulfur atoms of aminothiazole presented peaks at 165.15–172.25 ppm. The signal for C5 of the phenyl group of aldehyde appeared at 125.12 ppm for the ligand (S). Due to the presence of oxygen atoms, the carbon C14 of the phenyl ring of aminothiazole had key shifts, which were recorded at 153.92–153.93 ppm (more electronegative).[33] All the other carbon peaks were found to be in good agreement with the expected standards and cooperated fine with the entire number of carbon atoms currently in the projected constructions of the ligands.

Mass Spectra

The charge-to-mass (m/z) ratio was used to determine the structure from the mass spectra (Figures S13–S17). The ligand (S) spectra revealed a base peak of 100% intensity at 134 for the [C8H10O2•] fragment. This resulted in a peak at 405 (2.36%), with this being identical to its molecular weight (MW = 405) when the hydrogen radical was lost. Furthermore, this molecular ion was fragmented in two ways. The loss of the SH radical produced a fragment ion peak of 372 (2.36%), followed by the expulsion of the C10H6N2 molecule, which produced a fragment ion signal of 218 (2.36%) (14.98%). Further spectral compositions for the fragmentation of C=N, C–O, and C–C were seen at 235, 209, 199, 154, 116, and 105.[34] The most stable peak of the highest intensity for the ligand (S) was seen at 239 due to the fragment [C13H13N2OS]•. On subsequent expulsion of such a CO molecule, this fragmentation ion produced a fragment ion peak near 190 (100%), which is itself a base peak. On the loss of the NH2 radical, this base peak gives rise to a fragment ion signal at 174 (3.93%). The molecular ion fragmented in a different way, losing the SH radical and the C10H5N2 radical at the same time, yielding a fragment ion peak of 219 eV (2.36%). Following the elimination of the NCO radical, this fragment ion produced a fragment ion peak at 177 (38.58%). This ligand conceptual mass spectrum fragmentation trend was consistent with its structure, as shown in Scheme . Meanwhile, peaks for the remaining fragmentations of C=N, C–S, C–N, C–C, and C–N were seen at 339, 329, 291, 267, 252, 156, and 149.

Electronic Spectra

The stereochemistry and geometrical layout of metal complexes were assigned using molecular electronic spectrum measurements based on the sites and degree of d–d transitions. The electronic spectra of all the ligands and complexes were acquired in the DMF solvent at 298 K in the wavelength region of 200–800 nm. The electronic spectra of the synthesized ligands revealed strong bands at 253–298 nm due to an apparent π → π* transition inside the aromatic ring.[35]

In free ligands, the band identified at 306–392 nm was attributed to the azomethine unit with n → π* electronic interactions. The two bands attributed to the π → π* and n → π* transitions were moved to higher frequencies in metal complexes, which were attributable to metal chelation.[36] In addition to these two bands, d–d electronic transitions were also demonstrated. The charge transfer between aminothiazole from ligands and the metal ions was ascribed to the high-intensity absorption band of the metal complexes at max = 310–380 nm. All metal complexes exhibited an octahedral geometry based on evidence from electronic absorption experiments.

Molar Conductivity and Magnetic Studies

The molar conductance results for the metal complexes were documented at room temperature in a 10–3 M solution in the DMF solvent. The molar conductance values, ranging from 11 to 18 Ω–1 cm2 mol–1, indicated their non-electrolytic character.[37] The values of magnetic moment aid in forecasting the paramagnetic and diamagnetic behavior of metal complexes by providing critical information on the availability of unpaired electrons in metal ion d-orbitals. This investigation also aids in establishing the metal complex geometries. The magnetic moments of the cobalt complexes ranged from 3.82 to 3.91 BM,[38] supporting the octahedral geometry and paramagnetic character due to three unpaired electrons in the d-orbital. The nickel complexes had a value of 4.29–4.31 BM,[39] whereas the copper complexes had a magnetic moment value of 1.78–1.81 BM.[40] Such magnetic moment measurements justified the octahedral geometries assigned to the Ni and Cu complexes with two and one unpaired electrons correspondingly, demonstrating their paramagnetic characteristics. The zinc complexes had a zero magnetic moment value, confirming their diamagnetic character with octahedral geometry having no unpaired electrons.

Computational Details

Geometry Optimization

All of the optimized geometrical structures of synthetic compounds, at their ground-state energies, were considered using Gaussian 09 software and the density functional theory (DFT) methodology. The B3LYP approach with LanL2DZ basis sets was utilized to optimize geometries such that ligands and complexes might have the minimum energy-feasible structures[41] (Figure ). The benzene (C–C) system conceded DFT-optimized bonding distances of 1.41–1.42, 1.32–1.33, 1.33–1.35, 1.37–1.38, 1.32–1.34, and 1.32–1.33 Å. Bond angles of 120.0–121.3° are discovered throughout this process.[42] The expanded lengths appropriate for bonding of (C7–N14) and (N1–C7) were found to be 1.27 and 1.29 Å, respectively. In the context of their (C9–N10), (C11–N13), (C9–N10), and (C10–N11) bonds, the ligands were found to have (C9–N10), (C11–N13), (C9-N10), and (C10–N11) bonds. Their bond angles of 120.0–121.3° were discovered throughout this process. The ligands have been found to have increased lengths suitable for (C7–N14) (1.27 Å) and (N1–C7) (1.29 Å) bonding in the context of their (C9–N10), (C11–N13), (C9–N10), and (C10–N11) bonds (Figure S18).

Figure 1

Optimized geometries of ligands and metal complexes at the B3LYP level.

After the coordination with all the utilized metals with ligands,[43] the prominence of azomethine as a double bond (C=N) was lowered as a result of this scenario, which aided in chelation and impacted M–N bond development (Figure ). All the complexes had the twisted octahedral configuration with the central metal atom acting as a neutral bidentate ligand with carbonyl oxygen bound (M–N = 2.415–2.146 Å) and phenyl oxygen bound (M–O = 2.45–2.47 Å).

As the ligand had two nitrogen atoms in the molecular plane, their bonding as M–N = 2.53–2.54 Å and M–N = 2.48–2.51 Å was maintained in their octahedral geometry. The rigidity of ligands impacted the consistency of the metal complex geometries, causing the bond angles to diverge from 90°. The conventional charge of metals was also supposed to be lowered by the electron density donation by the active centers (Figure S19).

Frontier Molecular Orbital (FMO) Analysis

The analysis of frontier molecular orbitals (FMOs) may easily determine a chemical system with its optoelectronic capabilities as well as its capacity to absorb light. The energy differential among the highest occupied and lowest unoccupied molecular orbitals (HOMO and LUMO, respectively) was closely related to a molecule for its chemical reactivity and kinetic stability.[42] The HOMOs, which are electronically filled orbitals, have the option to contribute electrons, whereas the LUMOs, which are electronically vacant or unoccupied orbitals, have the potential to receive electrons. The modest difference in energy between HOMO and LUMO is relevant to the ability for a molecular design to polarize and also the possibility if there is intramolecular charge transfer for ligands to metals or vice versa.

The energies of frontier molecular orbitals (FMOs), mainly EHOMO-1, EHOMO, ELUMO, and ELUMO+1, and their energy gaps like EHOMO–LUMO and EHOMO-1–LUMO+1 remained the key characteristics to investigate the electrical characteristics of materials. These values were determined using B3LYP/6-311+G(d,p) basis sets. The antioxidant capacity and bioactivities of ligands are also directly connected to these values, exposing their most likely places that can be targeted by reactive agents and oxidants. The FMOs of ligands and reference compounds, in present study, were highly overlapping, indicating that the ligands are strongly reactive to antibacterial and antioxidant actions. The overall order for Egap’s was found as follows: (S) (2.68) > (S) (2.67) > terbinafine (2.51) > cefixime (1.72). The compounds with lower EHOMO values exhibited a weaker electron-donating aptitude, indicating that ligands may have stronger electron-donating capability than reference compounds, implying that these ligands would be better antioxidant and biologically active candidates. However, both the reference compounds had the lower Egap’s, which is also evident from their higher experimental activities. The satisfactory finding, in the present study, is indeed the comparable energy gaps of ligands and reference compounds.

In the ligand (S), the spatial extent of HOMO-1 was on the benzene ring as well as the oxygen of phenyl-hydroxyl, HOMO at the complete molecule, LUMO at the carboxylic group and the benzene ring, and LUMO at the hydroxyl groups. Intramolecular charge transfer (ICT) occurred between benzene and hydroxyl entities (HOMO and HOMO-1) and carboxylic and hydroxyl groups (LUMO and LUMO+1). In the ligand (S), the electron density mostly (HOMO-1 and HOMO) were on the thiazole moiety, LUMO at the benzoyl ring, and LUMO+1 at the benzoyl association. ICT can be noticed from the phenyl moiety (HOMO-1 and HOMO) to the thiazole circle (LUMO and LUMO+1) (Figure ). In cefixime (marketed antibacterial drug), the HOMO charge was mainly focused on its thiazole, which seemed quite different to that of the ligands. In its LUMO vicinity, however, the charge was dispersed on its main moiety. In terbinafine (marketed antifungal drug), the HOMO charge was mainly focused out of the benzene ring, while an opposite behavior was seen during its LUMO character. When ligands were contrasted to reference compounds, it was observed that they behaved similarly and that ligands may have greater electron-donating capability than reference compounds, hinting that these ligands might be attractive antioxidant and pharmacologically active alternatives.

Figure 2

Frontier molecular orbital investigation of ligands and reference drugs.

In metal complexes, intermolecular charge transfer (ICT) was ascertained from HOMO-1/LUMO, HOMO-1/LUMO+1, HOMO/LUMO, and HOMO/LUMO+1. All the Egap’s were found to be very narrow than the ligands and reference compounds to justify their enhanced bioactivities and antioxidant potentials. Their decreasing order was recorded as follows: (7) (5.94) > (8) (5.91) > (1) (5.62) > (6) (5.56) > (2) (5.51) > (5) (5.23) > (3) (5.18) > (4) (4.97). The extent of charge density of HOMO was established off the lone pair of nitrogen atoms, the methoxyphenyl and ethoxybenzyl components of the π-charge, and the M(II) ion of the dx2–y2 orbital. In ligand (S) metal complexes, ICT was also observed.

The charge density of HOMO and HOMO-1 was widely disseminated across the configuration, while the charge density of LUMO was widely dispersed throughout the structure except for the methoxybenzene portion, and the charge density of LUMO+1 was restricted to the thiazole moiety. The charge density of HOMO-1 and LUMO was exchanged on the partial structure, omitting the ethoxybenzene portion in the ligand (S) and metal complexes (5)–(8), with the charge density of HOMO mainly focused over the complete configuration and the charge density of LUMO+1 circulated on the thiazole moiety.

From the present investigation, it was assumed that there has been a donation of π-electrons, primarily from ligands, which contributes to a back-donation of the metal d-orbital as well as the d-orbital of metals toward the ligands, which jointly improves the metal complex stability (Figure ). The order of the overall energy gaps (Egap’s) and their subsequent reactivity was concluded to be as follows: reference compounds > metal complexes > ligands. The lesser energy gap attributes of complexes in contrast with the pure ligand, which forecasted good therapeutic efficacy, displayed the good attentiveness of the metal complexes. The nickel and cobalt compounds had the lowest value, implying that they are more effective inhibitors and this trend could be noticed from the experimental results.

Figure 3

Frontier molecular orbital analysis of DFT-optimized metal complexes.

Global Chemical Reactivity Descriptors

To recognize the association between the structural system, consistency, and global surface chemistry, interpretive density functional theory (DFT)-predicated global reactivity identifiers have been used. These identifiers have been used to create quantitative structure–activity (QSAR), structure property (QSPR), and structure toxicity (QSTR) relationships.

Electronegativity has been the most pertinent chemical attribute, which describes a species’ chance to acquire electrons to itself. The overall order for the synthesized and reference compounds was recorded as follows: terbinafine (6.77) > (S) (6.63) > (S) (6.44) > cefixime (6.13) > (7) (5.69) > (8) (5.53) > (6) (5.27) > (1) (5.14) > (2) (5.07) > (3) (5.02) > (5) (4.84) > (4) (4.61). The ligands and reference compounds had higher electronegativity values, which implied their higher reactivity as compared to the complexes. The aromaticity of a chemical is related to its (η). The (η) symbol characterizes the propensity of electrons to exit the electronic cloud. It also represents the degree of the electronic cloud conflict to deformation. The overall order for the synthesized and reference compounds was recorded as follows: (S) (1.34) > (S) (1.34) > terbinafine (1.26) > cefixime (0.86) > (1) (0.49) > (2) (0.44) > (5) (0.39) > (8) (0.39) > (4) (0.36) > (6) (0.29) > (7) (0.25) > (3) (0.16). The two ligands had higher electronegativity values, which implied their higher reactivity as compared to the complexes. The (ω) represents the system stability potential whenever it is saturated through electrons after the exterior conditions. The whole order for the synthesized and reference compounds was noted as follows: (3) (78.75) > (7) (64.75) > (6) (47.88) > (8) (39.64) > (5) (30.03) > (4) (29.52) > (2) (29.21) > (1) (27.18) > cefixime (21.85) > terbinafine (18.23) > (S) (16.4) > (S) (15.51). All the metal complexes had higher values than the ligands and reference compounds. The values of the Egap and reactivity descriptors showed that the ligands retain good reactivity. Any chemical arrangement with a minor FMO energy gap (Egap) is known for being less stable, more reactive, and softer.

The measurements of global hardness in the investigated ligands were higher than those of the measured results of metal complexes, indicating that the examined ligands were less reactive than their corresponding metal complexes. The hardness values and the reducing trend of FMO energy gaps were determined to be exactly matching. The chemical potential values have also been used to determine the reactivity and consistency of the compounds. Higher chemical potential numbers were thought to be less reactive and so more stable, and vice versa. Among all the synthesized compounds, the ligand (S) exhibited the highest value of ionization potential at 7.92 eV, whereas compound (3) exhibited the lowest value of ionization potential at 5.18 eV. The decreasing order of ionization potential was as follows: terbinafine > cefixime > (S) > (S) > (2) > (8) > (1) > (5) > (7) > (6) > (4) > (3). Because the ligands had the highest values of electron affinity than the metal complexes, the electron-acquiring and electron-donating capacity of the examined ligands was characterized and explained by the values of electron affinity and ionization potential, which corresponded with the energies of the HOMO and LUMO orbitals, respectively. In general, the ionization potentials were greater than their electron affinities, indicating that the examined ligands have the exceptional electron-donating capability, which also corroborated the results of the global electrophilicity index (ω). The ligand (S) also possessed the highest value of electronegativity at 6.01 eV, whereas compound (3) possessed the lowest value of electronegativity at 4.49 eV. Its declining order was recorded as follows: cefixime > terbinafine > (S) > (S) > (2) > (8) > (1) > (5) > (7) > (6) > (4) > (3). Based on the findings acquired, it was estimated that all the ligands were capable of behaving as chemically hard in their systems with superior kinetic stability and effective electrophilic capability (Table S3).

One-Electron Transferral Insights

The chemicals with so minimal EHOMO attributes have weak electron-contributory abilities, indicating that these compounds would be preferable for their electron-donating abilities. As a consequence, they may be capable of contributing to their improved antioxidant and physiological abilities. Natural antioxidants give a free radical an electron, forming a radical cation. Such a radical would have had to be stable enough with a single transfer framework to maintain radical scavenging capabilities. Ionization potential (IP) can be used to assess the antioxidant capacity, revealing the electron transfer context that can then be anticipated when IP = −EHOMO. For substances with relatively small IPs, the radical scavenging aspect is expected to be entertaining. These findings also suggest that ligands may have some antioxidant properties, which could help with measuring the radical scavenging efficacy.

Molecular Electrostatic Potential (MEP) Analysis

Any chemical system with its physical and chemical characteristics may be investigated using molecular electrostatic potential (MEP) plots. Such maps can be used to understand the potential for nucleophiles or electrophiles to engage at more suitable locations in chemical species.[42] The MEP surfaces have distinct hues such as green, orange, blue, red, and yellow that show the magnitudes of electrostatic potential within the chemical structures. The colors in their ranges as red > orange > yellow > green > blue were discovered to be the ascending order of magnitude of electrostatic potential. The oxygen atoms provided the red-highlighted region of the MEP map, which displays the area of negative potential and may be the most applicable. In contrast, the emphasized blue or green region indicates the area of positive potential and maybe the most appropriate place for nucleophilic engagement. The hydrogen atoms and certain carbon atoms that produce electron-deficient areas were usually represented by the colors blue and green, respectively. The negative and positive electrostatic predictors of ligands and reference compounds, respectively, were centered over the hydrogen and oxygen atoms of both the hydroxyl and carboxylic positions, as shown in the figure (Figure ).

Figure 4

MEP analysis of the ligand (S) and compound (5).

In the current work, negative electrostatic potential is seen on the oxygen atoms of ligands, whereas positive electrostatic potential is concentrated mainly on nitrogen and sulfur atoms. The investigation also revealed that the metal complexes had a high positive electrostatic potential dispersed across their skeleton, implying that they might be tightly attached to the microorganism under the current study. This conclusion could be useful in developing a stabilizing compound for optimal drug docking within the analyzed microorganism to a ligand binding scheme. The quantum chemical priorities would describe in full the biological exploit of these MEP surfaces in the molecules (Figure S20). This meant that these atoms were active places for cooperation with metal ions, which could bind with positive electron density to the components of the cell arrangement, etc.

NBO Numbering Scheme

This strategy is founded on the impression that it contributes evidence on intra/intermolecular interactions as well as bond similarities.[44] For chemical interpretations, this is an absolutely excellent method for exploring the hyperconjugative ability to participate and electron transmission by a charged lone pair of electrons.

The stabilization energy (E2) stands proportional to the interactions of the NBO intensities. The DFT/B3LYP/6-31G+(d,p) technique was used to calculate the natural bond orbitals, and accurate characterization of the information acquired by the second-order petulance theory investigation was evaluated (Table S4). This also provides a straightforward context for investigating boundaries in either full or empty orbital regions, as well as charge transformation and conjugation encounters in the chemical structure. This has been a worthwhile tool for determining the transferring/conjugation of charges in particular parts of a molecule organization with greater accuracy. The adoption of a second-order stabilizing capability can also explain the conjugation of the entire system. Nucleophilic hyperconjugation correlations are most likely to lead in stabilization frequencies of 2.93–3.19, 1.60–2.13, 4.03, 4.07 (S), and 1.18 (S) for transitions (Figure S21). The aggregation of the stabilization charges was inversely related to the occurrence of a hyperconjugative interface amid electron donors/acceptors. The back bonding appeared significant in all of these compounds, as large as the ligands or even larger than in the corresponding metal complexes. This was due to the strong σ-donor character, which renders typical metal acceptor sites particularly electron-rich and hence prone to back bonding.

Computed FT-IR Spectra

In experimental FT-IR, the number of atoms having either a symmetry point group or effectual vibratory normal patterns was determined to be the same (Figure ). The O–H bonds had a vibrational frequency of 3897, which is also given to it. The O–H absorption range occurred at 3700–3550 cm–1. For the previously mentioned wavelengths, the predicted modes were 3690, 3534, 3487, 3497, 3462, 3398, and 3387 cm–1 for molecules, which are very close to the experimental range. The absorption slopes for C–C oscillations were explored at 1640–1415 cm–1. The compounds with computed C–C bending frequencies in the aromatic rings were evaluated at 1650–1640 cm–1, which was very close to experimental measurements, attributed in the 1610–1520 cm–1 range. In complex (4), modeled C–C absorbance oscillations were ascertained at 1650–1635 cm–1, which aligned the precise results quite well (1555–1660 cm–1) (Table S5). They are made up of multiple tiny bands caused by the stretching vibration of the C–H alliance. In FT-IR, the aromatic ring C–H stretching frequencies for the current investigation are seen at 3231–3176 cm–1.[38] The experimental FT-IR spectra have been used to build the analysis for all of these excitation energies. The C–H oscillations of heteroatomic organic compounds and their metal complexes are extremely comparable to the benzene system C–H vibrations. In this exploration, a strong C–C vibration pattern was observed at 1721–1486 cm–1. The C–H vibrations were also investigated in the 1247–1087 cm–1 region. The C–H stretching strategies and their span are often confined to a lower frequency in recent studies.

Figure 5

TD electronic spectra of ligands and metal complexes.

The scales of the rocking stretch were 1326, 1296, and 1227 cm–1. Throughout heteroaromatic or aromatic twisting, the C–H absorption spectrum bands were quantified at 3100–3000 cm–1. The expected symmetric absorptions are 3170, 3202, 3012, 3182, 3140, 3141, and 3210 cm–1. Investigational vibrational wavelengths at 3153, 3210, 3195, 3155, 3210, and 3150 cm–1 are strikingly comparable to such vibrations. The extent of bending band C–N vibrations was a complicated task due to the risk of interaction with various other vibrations. It was noticed that the C–N stretch modes comprised 1600, 1445, 1286, and 1273 cm–1 (Figure S23).

TD-DFT UV–Vis Spectra

The CAM-B3LYP feature was used to perform the UV–Vis studies of the title molecules at their optimal geometries as part of the TD-DFT extents, at the equivalent basis set of (6-31G)+(d,p).[45] A comparative assessment likening investigational and computed (Figure ) study results was used to probe the molecular orbital representations, oscillator strengths (f), absorption wavelengths (λ), and vibrational energies, as seen in the constituent spectra. The theoretical UV–Vis spectra with absorption bands, the excitation, and oscillator strengths for the title compounds showed a band based at 1669–1656 nm with the oscillator frequency f1/4 = 0.0028, with the main H-2 → L (93%) contribution of charge transfer. With the oscillator amplitude f1/4 = 0.0021, the maximum absorbance strand was assessed at 1755 nm, with the transition being major contributions (63%) made by H → L (33%) and H → L + 1 (56%) (Figures S24 and S25). In contrast, the test findings only revealed two bands in their testing results: an intense, broad band centered at about 415 nm and also another at 390 nm.

Figure 6

Computed and experimental UV–Vis spectra of (a) ligands and (b) metal complexes.

Three bands throughout the electronic spectra of Cu(II) complexes (3) and (7), identified at 331, 366, and 437 nm, were illustrated by the π → π*, n → π*, and LMCT transformations. Zn(II) complexes (4) and (8) exhibited two bands at 350 and 435 nm, which were distinguished by the shifts such as π → π* and n → π*. A complete description of experimental absorption peaks, theoretical electronic processes, and their character designations can be found in Table S5 in the Supporting Information.

Natural Population Analysis

Atomic charges with molecules can be represented in a variety of ways, with a high level of authenticity to about the same values only after acknowledging the residual uncertainty, which is sufficient for molecular calculations. The Mulliken charges were derived through the measurement of Mullikan populations and are used to quantify partial atomic charges as an alternative to analytical chemical procedures.[46] These charges were determined by means of the DFT/B3LYP technique plus 6-31G+(d,p) as the basis set. The electron density of a molecular system on a wave function can be evaluated using the natural population analysis (NPA). The NPA localizes the electrons in each atom using natural atomic orbitals (NAOs) and natural bond orbitals (NBOs) with the highest possible electron density, reducing the dependence on the basis set to construct the natural charges from each atom. The far more positive or most negative NPA values are then markers of the electronic density distribution in the molecule as well as the possibility of drawing or donating electrons for the formation of covalent or noncovalent bonds. As a result, the spin density calculated using NPA can be utilized to explain the electronic dispersion in molecular systems (Figure ).

Figure 7

Mulliken charge analysis of the ligand (S).

A fair amount of positive or negative NPA values are predictors of the molecular electronic density distribution as well as the prospect of drawing or donating electrons for covalent or noncovalent bond formation. As a result, the electronic dispersion in molecular systems can be explained using the spin density calculated using NPA. At the B3LYP level of DFT, the atomic potential variation on the acceptor atoms of the investigated compounds donor atoms was approximated. This approach was used to examine the charge variations for both free ligand donor atoms and complex acceptor atoms. So, when atomic intensities of the donor atoms and the M(II) cation electrons were obtained, it was evident that the donor atom concentration in the complexes was substantially lower than that in the free ligand, while the M(II) cation electron intensity increased (Figure S26).

Antibacterial Activity

The antibacterial screening was performed on the ligands (S) and (S), as well as their metal complexes (1)–(8), toward two bacterial strains: Micrococcus luteus (Gram-positive) and Escherichia coli (Gram-negative) (Table ). M. luteus belongs to the Micrococcaceae family and is a Gram-positive to Gram-negative, nonmotile coccus, having a tetrad shape, and a pigmented and saprotrophic type of bacterium. Various E. coli classes are mostly harmful; however, rare classes can affect the host considerably by food poisoning. Correspondingly, sometimes, they might be liable for stomach problems. The overall order of reactivity for all the compounds against the two bacterial strains was recorded as follows: cefixime > azithromycin > (8) > (4) = (5) > (3) > (6) = (7) > (2) > (S) > (S) = (1) (M. luteus) and cefixime > azithromycin > (8) > (7) > (S) = (5) > (4) > (6) > (S) = (1) = (2) > (3) (Figure ).

Table 1

Antimicrobial Activity of Ligands and Metal Complexesa

	bacterial species			fungal species
comp.	M. luteus	E. coli	SA	A. niger	A. terreus	SA
(S1)	13	12	12.50	15	14	14.50
(S2)	12	15	13.50	14	12	13.00
(1)	12	12	12.00	13	14	13.50
(2)	14	12	13.00	14	15	14.50
(3)	16	11	13.50	17	15	16.00
(4)	17	14	15.50	18	14	16.00
(5)	17	15	16.00	16	14	15.00
(6)	15	14	14.50	19	18	18.50
(7)	15	21	18.00	20	17	18.50
(8)	20	19	19.50	18	16	17.00
SD1	31	30	30.50
SD2	32	31	31.50
SD3				28	27	27.5

SD1: azithromycin; SD2: cefixime; SD3: terbinafine.

Figure 8

Antibacterial activity of ligands and metal complexes.

Their antimicrobial influence was also compared to that of azithromycin and cefixime, the two standard drugs used to inhibit the bacterial strains M. luteus and E. coli in 32 and 30 mm zones, respectively. The ligand (S) was moderately active against E. coli, while a maximum activity of 15 mm toward M. luteus was found. The ligand (S) was determined to be the most active ligand among the synthesized ligands, with the highest activity of 15 mm. Compound (8) had the maximum activity of 21 mm against E. coli and moderate activity (11 mm) against M. luteus among these derived metal complexes for compound (1). However, complex (3) was the least active, displaying inhibitory action against E. coli and Salmonella typhimurium.

In comparison to both the standard drugs, the antibacterial data revealed that most of the investigated compounds had moderate to substantial antibacterial activity (standard drug).[47] The metal complexes, on the other hand, have a stronger antibacterial activity over their uncomplexed ligands. The antibacterial findings showed that ligands and metal complexes played a substantial role in the bactericidal activity of the compounds. It was clear that the chelation had increased the bioactivity of the ligands. On the basis of chelation theory, complexes with larger zones of inhibition, unlike ligands, could be perceived. Because of the engagement between the ligand orbital and the composition of the positive charge of metallic ions associated to donor groups, polarization for metal ions is minimized to a significant degree through chelation. It improves electron polarization around the chelating ring, as well as complex permeation into lipid membranes and subsequent hindering of metal adsorption sites in microorganisms.

Antifungal Activity

The antifungal results were then compared to that of terbinafine, a commonly used antifungal drug. The data revealed that the majority of the compounds exhibited significant antifungal (12–19 mm) efficacy against the fungal strains (Table ). Aspergillus niger is a prevalent food contaminant that allows a disease known as “black mold” to certain fruits and vegetables. It is found all over the world in soil and is frequently reported from indoor spaces. Aspergillus terreus, also called Aspergillus terrestris, is a fungus that can be found in soil all over the world. The warmer climate conditions, such as tropical and temperate areas, are home to this saprotrophic fungus.

Table 2

Antioxidant Activity for Synthetic Ligands and Their Metal Complexes

	DPPH		ferric reducing power
comp.	inhibition (%)	IC50 (μL)	inhibition (%)	IC50 (μL)
(S1)	52.0 ± 0.11	104 ± 0.11	54.3 ± 0.11	108 ± 0.08
(S2)	70.0 ± 0.13	146 ± 0.12	58.3 ± 0.11	116 ± 0.09
(1)	58.2 ± 0.10	116 ± 0.09	61.3 ± 0.12	122 ± 0.10
(2)	60.0 ± 0.09	120 ± 0.10	55.1 ± 0.10	110 ± 0.07
(3)	72.0 ± 0.11	144 ± 0.11	62.3 ± 0.11	124 ± 0.12
(4)	63.0 ± 0.12	126 ± 0.12	56.1 ± 0.13	112 ± 0.13
(5)	54.5 ± 0.10	108 ± 0.13	57.3 ± 0.11	114 ± 0.09
(6)	67.0 ± 0.11	134 ± 0.10	59.1 ± 0.12	118 ± 0.10
(7)	62.0 ± 0.12	124 ± 0.11	68.1 ± 0.10	136 ± 0.12
(8)	71.2 ± 0.11	142 ± 0.12	66.3 ± 0.09	132 ± 0.11

It has been located in a range of habitats, including decaying vegetation and debris, in addition to soil. Common synthetic acids and also enzymes like xylanase are produced by A. terreus in industry. It was also the first source of the drug mevinolin (lovastatin), which lowers serum cholesterol. The ligands demonstrated a modest antifungal action (12–15 mm). The ligand (S) demonstrated the highest activity (15 mm) against A. terreus and the lowest (14 mm) against A. niger. Similarly, the ligand (S) had moderate (12–14 mm) action. All of the metal complexes (1)–(8) validated moderate to good activity for the two fungal strains (Figure S32). The A. terreus strain inhibited compound (4) to the greatest extent (29 mm), while the same strain inhibited compound (8) to the least extent (19 mm). It was revealed that chelation with metals increased the antifungal action of ligands. Compound (7) showed the highest (28 mm) activity against A. niger of all the results, whereas compound (8) had the lowest (12 mm) activity against the same species (Figure S26). The current research has also established that chelation, in the circumstance of antifungal action, continued to generate a more complete and constant biocidal influence, inhibiting the formation of bacteria, which is compatible with the study findings. New studies has also demonstrated that in the sense of antifungal activity, chelation continues to deliver a more consistent and accurate fungicidal aspect, inhibiting the propagation of microorganisms, analogous to the findings of this study.

For antimicrobial investigations, the metal complexes (1)–(8) were shown to be significantly bioactive than the ligands, and this action was enhanced by coordination with metal ions.[48] However, when compared to standard drugs, the complexes were only moderately active. The fact that the ligands had a double carbon-nitrogen link may have contributed to the increased activity of the complexes. It was discovered that the metal ions in the complexes possessed a positive charge, implying that it is somewhat complementary to the donor atoms in ligands and that π-electron delocalization may occur throughout the chelate.[49] The metal complexes gained a lipophilic feature as a result of the chelation, which helps their absorption through the lipid membrane of microbes.[50] The solubilization, contact, permeability, and length of the link between the metal and ligand lipid layers of bacterium membranes are all elements to consider.[51]

Antioxidant Activity

The definition of antioxidant activity is “the inhibition of the rate of increase in oxidative chain reactions to restrict the oxidation of proteins, fatty acids, DNA, or other substances”. Secondary antioxidants indirectly prevent free radical formation via the Fenton reactivity approach, while primary antioxidants consciously harvest free radicals. Diphenyl picryl hydrazide (DPPH), being a stable free radical, is used to check the radical scavenging ability.

The results presented for DPPH (%) activity showed that complex (3) had the highest (72.0 ± 0.11%, IC50 = 144 ± 0.11 μL) antioxidant activity, while the ligand (S) exhibited the lowest (52.0 ± 0.11, IC50 = 104 ± 0.11 μL) antioxidant activity. The overall order of antioxidant activities was recorded as follows: (S) > (3) > (8) > (6) > (4) > (7) > (2) > (1) > (5) > (S). The ferric reducing power works on the principle of arrangements with reduction potential reacting with ferric chloride (Fe3+) to create potassium ferrocyanide (Fe2+), which then reacts with ferric chloride to form a ferric–ferrous complex with a maximum absorption at 700 nm. With the increase in supply, the strength and efficiency of hydro alcoholic extracts decrease. The total iron reducing power, measured in milligrams of gallic acid equivalence per gram (mg GAE/g), was found to be in excellent condition. Compound (7) showed the maximum (68.1 ± 0.10%, 136 ± 0.12 μL) activity, while the ligand (S) was concluded as the least active (54.3 ± 0.11%, 108 ± 0.08 μL) compound as it presented the least activity in these sample concentrations (Figure ): (7) > (8) > (3) > (1) > (6) > (S) > (5) > (4) > (2) > (S).

Figure 9

Antioxidant activity of ligands and metal complexes.

According to a literature review, antioxidant actions and the total phenolic content have a very synergistic relationship, which was also evident in our research designs (Figure S28). The results of the quantitative assessment of antioxidant activities revealed a positive correlation between such calculations, implying that many entities play a significant role in being a good antioxidant,[52] which was also visible in our correlational analyses. The quantitative examination of antioxidant activities revealed a strong association for each other.

Druglikeness Study

In silico druglikeness assessment is used to see if hypothetical compounds or molecules have a pharmacokinetic profile that can be evaluated to compare with their synthesized versions. The druglikeness attributes of the compounds with appropriate physicochemical characteristics were determined by applying one of several filtering rules, namely, Lipinski’s rule, by a computational engine called Molinspiration (version 2018). Because most of the compounds, the ligands and metal complexes (1)–(4), had a molecular weight of 500, it was possible to predict that such complexes would be easy to transport (Table S7). The log p (octanole water partition coefficient) (−6.46 to 2.06), which also generally quantified biochemical lipophilicity (oral availability), and TPSA (total polar surface area), which would be associated with a molecule’s hydrogen bonding and would be a great predictor of the drug molecule with its bioavailability, are both within acceptable limits (59.18–54.72). As a result, all hypothetical compounds met the criteria for an orally effective dose and could be established further as oral active compounds. All the other related properties are summarized in Figure .

Figure 10

Druglikeness analysis of ligands and metal complexes by Lipinski’s rule.

Conclusions

Two new thiazole-based ligands were synthesized and studied utilizing physical, spectral, analytical, and magnetic methods in this study. The octahedral geometries were proposed for respective bivalent metal (Co2+, Ni2+, Cu2+, and Zn2+) complexes established on spectral and magnetic study findings. Their non-electrolytic nature was proved by their molar conductivity. At the B3LYP/level of theory, quantum chemical computations based on DFT/TD-DFT studies were successfully applied to achieve optimized molecule structures with electronic calculations. The computed values of global hardness were lower than the intended values of global softness, implying that the examined ligands were more reactive and less stable. Antimicrobial and antioxidant tests were performed on all of the produced compounds. The antimicrobial and antioxidant activities of all of the compounds were comparable to those of standard drugs; however, cobalt and zinc complexes showed substantial antimicrobial property, while Ni(II) and Cu(II) complexes had more prominent antioxidant capabilities. The heterocyclic (thiazole) ring comprising the N and S heteroatoms was responsible for the massive outcomes. The chelation resulted in metal complexes with greater biological activity than their respective ligands. All of the novel compounds were determined to be stable and might potentially be employed in the development of orally administered candidates. Because of their antimicrobial properties, the synthesized chemicals might aid the pharmaceutical sector in reducing or inhibiting pathogen development. This research revealed that multidrug resistance can be addressed in the future by developing metal-based medicines based on a variety of newly developed pharmacological specifications.

Experimental Section

Materials and Measurements

The entire investigation was done with Sigma-Aldrich analytical-grade chemicals, salts, and solvents (St. Louis, Missouri, United States). Ethanol was distilled two times before being consumed. In the Supporting Information, there is information about the instruments. On a Stuart melting point instrument, the melting points of synthetic samples were confirmed. The FT-IR spectra were detailed in their spectral zones that used a Nicolet spectrophotometer within the default setting. A Bruker Avance 300 MHz instrument was used to determine the 1H and 13C NMR spectra. Synthesized derivatives (ligands and metal complexes) were prepared by dissolving 10 mg and pumped directly through into an LTQ XL Linear Ion Trap Mass Spectrometer (Thermo Scientific, USA) installed including an electrospray ionization (ESI) sensor for spectrometric assessment. Their UV–visible measurements were performed using a high-tech Shimadzu UV-4000 spectrophotometer. The Natural Product Research Laboratory, University of Gujrat, Gujrat, evaluated biological activities using the recommended procedures.

Synthesis of Ligands

The ligands were developed by expending the previously described method.[43] The ligand (S) was designed by adding a magnetically stirred, 10 mmol ethanolic solution (1.52 g) of 2-hydroxy-3-methoxybenzaldehyde to a 10 mmol ethanolic solution of 2-methoxy-6-((thiazol-2-ylimino)methyl)phenol. The mixture was refluxed for 3 h at 100 °C, and the TLC method was used to monitor the progress of the reaction. The response was considered complete when the color changed and a single TLC spot appeared. The solvent was evaporated over a rotary vacuum evaporator to recover the product. Washing the product in a hot ethanol/ether (1:1) solution further purified it. The ligand (S) was designed in the same way, except that the same aldehyde was refluxed with 6-ethoxy-1,3-benzothiazole-2-amine instead of the same aldehyde.

2-Methoxy-6-{[(1,3-thiazol-2-yl)imino]methyl}phenol, C11H10N2O2S (S)

MW: 234.17.2; yield (%): 73; mp (°C): 98–100; color: brown-red; FT-IR (cm–1): 3434 (OH), 2974 (OCH3), 1634 (HC=N), 1417 (C–N), 1345 (C–O); 1H NMR (d ppm): 3.43 (s, 3H, OCH3), 6.95 (d, C4-H), 6.98 (t, C5-H), 7.63 (d, C6-H), 8.27 (s,1H, HC=N), 11.207 (s, 1H, OH) (d, 2H-thiazole); 13C NMR (δ ppm): 56.28 (methoxy), 111.11 (C2, thiazole), 139.28 (C3, thiazole), 150.60 (C1), 139.88 (C2), 118.09 (C3), 129.33 (C4), 123.10 (C5), 111.14 (C6), 163.97 (C1-thiazole); MS (ESI) m/z: 361.37, 259.28, 226.25, 182.19, 150; anal. calc. (%): C (56.39), H (4.30), N (11.96); found: C (56.35), H (4.27), N (11.94).

2-{[(6-Ethoxy-1,3-benzothiazol-2-yl)imino]methyl}-6-methoxyphenol, C17H16N2O3S (S)

MW: 328.39; yield: 68%; mp (°C): 145–146; color: yellow-orange; FT-IR (cm–1): 3267 (OH), 2922 (OCH3), 1634 (HC=N), 1467 (C–N), 1347 (C–O); 1H NMR (δ ppm): 6.91 (d, C3-H), 7.23 (d, C4-H), 6.77 (s, C6-H), 7.55 (s, C15-H), 7.62 (d, C12-H), 7.59 (d, C13-H), 3.96 (q, CH2), 1.33 (t, CH3); 13C NMR (δ ppm): 56.39 (methoxy), 106.63 (C15), 113.82 (C2-thiazole), 114.25 (C6), 118.49 (C7), 119.63 (C3), 124.60 (C3-thiazole), 127.24 (C5), 132.29 (C4), 137.69 (C2), 183.81 (C8), 165.17 (C9), 147.14 (C10), 149.65 (C11), and 153.93 (C1 thiazole); MS (ESI) m/z: 312, 333, 353, and 373; anal. calcd. (%): C (62.18), H (4.91), N (8.53); found (328.39): C (62.18), H (4.91), N (8.53).

Synthesis of Metal Complexes

The metal complexes were synthesized, taking a metal-to-ligand ratio of 1:2, using the previously revealed method.[34] A 50 mmol solution of the appropriate metal salt in ethanol was added toward a magnetically refluxed (100 mmol) ethanolic solution of a ligand. The resulting mixture was refluxed for 6 to 8 h, resulting in product precipitation. To further purify the filtered product, it was rinsed in hot ethanol.

Antimicrobial Activity

The antimicrobial activity, evaluated as antibacterial and antifungal, for all the synthetic products was deliberated (in vitro) by the disc diffusion bioassay method.[38] During the investigation, one Gram-positive bacterial species (M. luteus) and one Gram-negative bacterial strain (E. coli) were tested, as well as two fungal strains (A. niger and A. terreus). The bacterial inoculation medium was developed by adding 2 g each of nutrient broth and agar before autoclaving it for antibacterial activity. The medium was cooled to room temperature before being mixed with a bacterial culture and poured into Petri plates. The sterilized discs were then placed on top of the settled mixture in Petri dishes. The micropipette pinched a 100 μL sample on the Petri plates. The resulting culture was incubated at 37 °C overnight to measure the bacterial inhibition (mm) zones. The standard drugs, azithromycin and cefixime, as well as the solvent (DMSO) in the same concentration (2 mg/mL), were used as controls.

The antifungal activity was also assessed by repeating the technique, except that terbinafine (an antifungal drug) was used as a reference drug when the inoculum was introduced into potato dextrose medium.

The antiradical (DPPH) scavenging (%) behavior and total iron reducing power (%) were used to calculate antioxidant activity. A quantitative connection between all the assessments was inferred by using each of the two antioxidant readings in the correlation diagrams in the sequence. The R2 correlation analysis was employed to assess the interdependence of such activities.

Antiradical Scavenging of DPPH

The antiradical scavenging of DPPH (%) was determined using the previously mentioned method.[53] In a Pyrex tube, a methanolic (0.1 mL) sample solution (1.0 mg/mL) was dipped and 4 mL of pure methanol was added, accompanied by 1.0 mL of 2,2-diphenyl-1-picrylhydrazine (DPPH). The generated mixture was permitted for 30 min at room temperature to evaluate the absorption of the resulting solution at 550 nm. To compare results, the reference standard butylated hydroxytoluene (BHT) was tested in combination with the sample. The inhibition (%) was calculated expending eq :

Ferric Reducing Power

The ferric reducing power (%) was calculated by using the previously mentioned route.[43] A test tube was sequentially admitted to 1.0 mL (1.0 mg/mL), 2.4 mL (2.0 M, pH 6.6), and 5.0 mL (10%) of potassium ferricyanide. For 30 min at 50 °C, the resulting mixture was incubated. After incubation, 5.0 mL of 10% trichloroacetic acid solution was introduced for almost 10 min to the centrifuge at 10,000g. The supernatant layer of 0.5 mL was isolated upon enabling the solution to be left for 30 min, and was assigned to measure its absorption at 700 nm with 1.0 mL of 1.0% ferric chloride. Butylated hydroxytoluene (BHT) was also used to relate the results of all samples as a reference sample solution. In their 3-fold measurement, the results were averaged.

Theoretical Studies

The Gaussian 09 D.0.1 package[54] was used to study DFT- and theoretical simulation-based molecular mechanics. The optimized structures, theoretical spectra, and summary of their geometries like bond angles and bond lengths were viewed using Chemcraft version 1.6, GaussView version 5.0.9, and GaussSum 3.0.2 software.

Optimization of Molecular Geometries

Using the B3LYP function of DFT and the LanL2DZ basis set, the molecular configurations of ligands and metal complexes were adjusted to the ground-state energies.[55] After optimization, the geometries did not exhibit hypothetical vibrational frequencies, confirming that they were low-energy structures. All calculations were performed with the isolated molecules in the gas phase. The CAM-B3LYP function was used in collaboration with the 6-311+G(d,p) basis set and the time-dependent density functional theory to do the theoretical UV–Vis assessment (TD-DFT). The highest molecular orbital energy (EHOMO), the lowest unoccupied molecular orbital energy (ELUMO), and associated energy differences (EHOMO–LUMO) for their optimal structures were found to infer their quantum chemical metrics. The theoretical FT-IR spectra were generated using DFT/6-311+G(d,p) functionalities at their optimum configurations.

Global Reactivity Identifiers

The energies of FMOs (Egap’s) were used to evaluate global reactivity identifiers including that kind of electron affinity (EA), ionization potential (IP), electronegativity (χ), global softness (σ), global hardness (η), global electrophilicity index (ω), and calculated chemical potential (μ) using formulas.

The global chemical reactivity descriptors (GCRDs) are imperative metrics for understanding the reactivity and structural permanency of any chemical system.[56] These descriptors are also known as biological activity descriptors. We evaluated the chemical hardness (η), chemical potential (μ), electronegativity (χ), global softness (σ), and electrophilicity index (ω) of the ligands and their metal complexes (eqs –8). The energies of HOMO and LUMO were operated to analyze these promising descriptors and are represented in Table S3.

The ionization potential is denoted by IP (eV), and the electron affinity is denoted by EA (eV).

Koopmans’s theorem was used for the determination of chemical potential (μ) and electronegativity (χ), in addition to chemical hardness (η), which are described and calculated from the following given equations as:

The following relationship was used to define the global softness (σ) as:

Meanwhile, the electrophilicity index (ω) can be explained and calculated by the given equation:

Natural Bond Orbital Analysis

The optimized geometry at the very same level with the B3LYP functional, respectively, was used to conduct natural bond orbital (NBO) interpretation with NBO Edition embedded in the Gaussian 09 suite for spin-paired as well as spin-unpaired structures. The stabilization energy (E2) of the delocalized donor (i) was determined for this purpose by applying eq to the acceptor (j) orbits.

In Silico Studies

Lipinski’s rule of five aids in the differentiation of druglike or non-druglike molecules. It forecasts a high likelihood of successes or failures due to drug similarity for molecules that meet two or more of its specific criteria. The Molinspiration engine (version 2018.10) was used to predict and display physicochemical characteristics in order to identify whether they are tied to synthesized samples in terms of bioactivity and hydrophobicity.