Ligational behavior of bidentate NO-donor, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol towards VO²⁺, Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺ Zn²⁺ and ions preparation, experimental, theoretical characterization, and microbicidal investigation

Abstract

Novel bidentate chelator, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol was employed for designing and preparation of transition metallic complexes. The synthetics were characterized based on various analytical, spectral, and theoretical tools. Based on the various analysis data, the azo-Schiff base chelator (H₂L) linked with metallic cations as a neutral/uni-negative bidentate chelator via the protonated/deprotonated hydroxyl oxygen and azomethine nitrogen atoms adopted a distorted octahedral, tetrahedral, or square planar geometry. The mass spectroscopy confirmed that the molecular weights of azo-Schiff base and some of its complexes closely match with that of the estimated m/z values. Additionally, the structure of the azo-Schiff base ligand and some metallic complexes were studied by Density functional theory (DFT) calculations to assess the molecular orbitals (LUMO&HOMO) as well as the molecules electrostatic potential (MEP) to complete the picture. The synthetics were evaluated for their in vitro activity against Escherichia coli, Aspergillus niger, and Bacillus subtilis. According to this research, unbound ligand molecules are inactive while their metallic complexes have low to moderate activity.

KEYWORDS:

• Azo-Schiff base
• DFT
• metallic complexes
• microbicidal activity

1. Introduction

One of the treasures of discovering new chemical compounds with biological and pharmacological activity is coordination chemistry. It is a brand-new, highly beneficial source of synthetic and semi-synthetic pharmaceuticals that enhances the activity of drugs to overcome the pathogens resistant [1]. The azo-Schiff base ligand was created when azo-dyes moiety (-N=N-) and Schiff-base with functional azomethine linkage (-C=N-) are combined [2]. The physiochemical and biological features of this flexibility class of chelators and their associated chelated complexes are improved by the inclusion of two active moieties in a single molecule [3]. The therapeutic properties of azo-compounds was discovered in Prontosil, sulphonamide associated azo-dye, which used as a medication against microbial infections [4,5]. The biological properties of azo-compounds have made them a focus of attention due to their potential as fungicidal [6–9], anti-oxidant [10], anticancer [11], anti-inflammatory [6,12], antiviral (HIV) [13], and bacteriostatic agents [6,7]. The Schiff bases with distinctive azomethine group were also aroused the attention of many researchers because of their biological activities and excellent therapeutic applications as well as their chelating properties [14]. It has previously been reported that the presence of the azomethine moiety(-C=N-) is primarily responsible for their biological activity [15–17]. Consequently, compounds with an azomethine moiety represent a significant class of compounds as pharmaceutical and therapeutic agents with a broad variety of pharmacological effects with analgesic [15], antiproliferative, anticancer [18–22], anti-inflammatory [15,23], antiepileptic [24], antitumor [25–28], antibacterial [29–33], antidepressant [34], antifungal [30–33], anti-oxidant [19,25–27], antidiabetic [35,36], enzyme & protein inhibitor [23] herbicidal [27,37], and antiviral viruses (HIV, influenza A, COVID-19 and a21, denovirus) [26,27,38]. Deswal et al. showed that the Co(II), Ni(II), Cu(II) and Zn(II) complexes of 2-(((3-mercapto-5-(pyridin-3-yl)-4H-1,2,4-triazol-4-yl)imino)methyl)-4-nitrophenol have moderate to good alpha-amylase and alpha-glucosidase inhibitory effect [39]. Nowadays, Schiff base metal complexes have employed as precursors and effective low-cost production techniques for preparing metal oxide nanoparticles by thermal decomposition [40]. Schiff bases are also an significant class of compounds that are highly instrumental in formation of complexes as well as act as sensors for various metal ions [41,42]. Recently, metal complexes of Schiff base assembled from natural products which used for coatings displayed antibacterial-renew dual-function anti-biofouling activity [43] and used as catalysis in the several reaction as their action as catalysts in reduction of nitro aromatic compounds [42], oxidation of primary and secondary alcohols [44] as well as in the Suzuki–Miyaura and Mizoroki-Heck cross-coupling reactions [45]. The current work studies the novel metallic complexes of azo-Schiff base incorporating both the azomethine (C=N-) and Azo- (C–N=N–C) linkages. It aimed to synthesis of VO²⁺, Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺ Zn²⁺ and ${\text{UO}}_2^2$ complexes of new azo-Schiff base namely, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol (H₂L, 1) (Scheme 1). Consequently, physio-chemical, and analytic techniques such NMR, UV-Vis, FT-IR, elemental analyzer, ESI-MS, molar conductance, magnetic moment, and thermal measurements as well as theoretical calculations were used to characterize the structure of the produced compounds. The geometry of the chelator and its metallic complexes was examined auxiliary to the density functional theory (DFT) calculations at the B3LYP/6-311++G(d, p) levels. To complete the picture, we assessed not only the overall energy but also the molecular orbitals (LUMO&HOMO) and the molecular electrostatic potential (MEP). Additionally, the in vitro antifungal and antibacterial activities of the chelator and its complexes was examined to discover novel, highly effective antimicrobial compounds to address the issue of antimicrobial resistance that most commonly occurs with the usage of medications.

Scheme 1. Scheme for the preparation of metal complexes in ethanol/Reflux 5 hs.

2. Materials and methods

The DMSO, absolute ethanol 1H-1,2,4-triazol-5-amine, and metal salts, (assay = 98.0–99.9%) were bought from Merck and employed with no extra purification. A published method was used to prepare 2-hydroxy-5-(o-tolyldiazenyl)benzaldehyde [46].

2.1. Apparatus

In Cairo University's Micro-Analytical Laboratory, Egypt, the elements (N, H, and C) for the produced compounds were analyzed, while the metal contents (VO²⁺, Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺, Zn²⁺) and Cl^- ions were assessed using conventional analytical techniques [47,48]. On an infrared spectrophotometer (Perkin–Elmer-1430) at National research centre, Egypt, KBr plates were employed to measure the IR spectra over of 400–4000 cm⁻¹ range. On a SHMADZU 2600 spectrophotometer, the electronic absorption spectrum was measured using 1-cm quartz cell over of 200–1100 nm range in DMSO at university of Jeddah. On mass spectrometer (JEUL JMS-AX-500), at National research centre, Egypt, the mass spectrometry was determined while on a JEOL EAC-400 spectrometer at Qassim University, Saudi Arabia, the NMR spectra was measured in DMSO-d₆. The complexes TG analysis was measured on a Perkin Elmer 7 thermal analyzer over a range of 25–800°C at a heating rate of 10 °C/min at National research centre, Egypt. At university of Jeddah, the magnetic susceptibility was recorded at 25°C using a Gouy Matthey’s Balance and computed by a published equation [49]. Tacussel type CD6NG conductivity bridge was used to test the complexes’ molar conductance (10⁻³ M/DMSO), and the published equation was used to compute it [50]. Unfortunately, there were countless failed attempts to get a single crystal. The clarity of prepared compounds has been verified by TLC utilizing various mobile phases.

2.2. Synthesis of azo-Schiff base ligand (H₂L)

The ligand, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol (H₂L, 1) was prepared by refluxing a mixture of 2-hydroxy-5-(o-tolyldiazenyl)benzaldehyde (240 mg, 1 mmol, 30 mL of EtOH) and 1H-1,2,4-triazol-5-amine (84 mg, 1 mmol, 50 mL EtOH) in presence of 0.1 mL glacial CH₃COOH for five hours. Once the coloured solid had formed, the reaction mixture was allowed to cool at 25°C before being separated, washed with EtOH, and dried in a vacuum to produce azo-Schiff base (H₂L, 1) (271.7 mg, 0.887 mmol, 88.7%) (Figure 1) as a yellowish green solid, (Melting point) m.p. = 240°C. Anal. (Elemental analysis) for C₁₆H₁₄N₆O (306.3 g /mol): Found (calcd.) %C 62.46 (62.74), %H 4.42 (4.61), %N 27.37 (27.44). FT-IR (KBr, cm⁻¹), 3398 ν(²²O-H³³), 3233 ν(³⁰N-H³²), 3080w, 3000, 2964w, 2888w ν(C–H), 1660 ν(²³C = N²⁴), 1605 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1475 ν(¹¹N = N¹²), 1267 ν(²⁰C-O²²), 1000 ν(²⁹N-N³¹). ¹H-NMR (400 MHz, DMSO-d₆): δ(ppm) =  14.37, 8.35 (s, 1H, NH³⁰), 12.619 (s, 1H, OH³³),7.96 (d, 2H, C-H^8&9), 8.395 (m, 1H, CH^7&10), 7.31 (d, 1H, CH²¹), 8.819 (d, 1H, CH¹⁷), 9.13 (s, H, CH¹⁹), 9.56 (s, H, CH²⁵), 8.173 (s, H, CH²⁸), 2.67(s, 3H, CH₃). ¹³C-NMR (90 Mhz, DMSO-d₆): δ(ppm) = 134.17 (C1&5), 122.42 (C2&4), 150.63 (C3), 140.42 (C6), 144.12 (C13), 128.29 (C14), 125.51 (C15), 113.79 (C16), 117.21 (C18), 166.31 (C20), 163.42 (C23), 159.58 (C26), 146.6 (C27), 27.45 (C34).

Figure 1. Syntheses of 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol (H₂L, 1).

2.3. General procedure for the synthesis of metal complexes

VO²⁺, Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺, Zn²⁺, and ${\text{UO}}_2^{2+}$ complexes were synthesized by refluxing an equimolar amounts of the following salts’ ethanolic solution (50 mL): VOSO₄·3H₂O (217.0 mg), FeCl₃·6H₂O (271.3 mg), RuCl₃·3H₂O (261.5 mg), Co(CH₃COO)₂·4H₂O (249.1 mg), Mn(CH₃COO)₂·4H₂O (245.1 mg), Ni(CH₃COO)₂·4H₂O (248.8 mg), Cu(CH₃COO)₂·2H₂O (199.7 mg), CuCl₂.2H₂O (170.5 mg), Cu(NO₃)₂.3H₂O (241.6 mg), CuSO₄·5H₂O (249.7 mg), Zn(CH₃COO)₂·2H₂O (219.5 mg), and UO₂(CH₃COO)₂·2H₂O (424.2 mg) with 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol (H₂L, 1) (306.3 mg, 30 mL of EtOH). For four hours, the refluxing was done while being stirred. After heating and rinsing many times with hot EtOH, the resulting solid-colored complexes separated out and were eventually dried under vacuum over P₄O₁₀.

2.3.1. VO²⁺-complex (2)

Color: light brown, yield (50.6%), mp = 266 ^oC, molar conductance (Λ_m) = 34.9 ohm⁻¹cm²mol⁻¹. Anal. for [VO(H₂L)SO₄(H₂O)₂], C₁₆H₁₈VSN₆O₈, (505.36 g/mol): Found(calcd.) %C 38.52(38.13), %H 3.31(3.59), %N 17.01(16.63), %V 9.77(10.08). FT-IR (KBr, cm⁻¹), 3488 ν(²²O-H³³), 3422 ν(H₂O), 3188 ν(³⁰N-H³²), 3046w, 2920, 2875w ν(C–H), 1633 ν(²³C = N²⁴), 1601 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1474 ν(¹¹N = N¹²), 1319 ν(²⁰C-O²²), 1022 ν(²⁹N-N³¹) 1155, 1061 ν(SO₄), 970 (O = V), 577 (²²O→V), 491 ν(²⁴N→V).

2.3.2. Fe³⁺-complex (3)

Color: reddish brown, yield (56.4%), mp = 280 ^oC, Λ_m= 22.2 ohm⁻¹cm²mol⁻¹. Anal. for [Fe(HL)Cl₂(H₂O)₂].2H₂O, C₁₆H₂₁FeN₆O₅Cl₂, (504.13 g/mol): Found (calcd.) %C 38.09(38.12), %H 3.98(4.20), %N 16.47(16.67), %Cl 13.87 (14.06), %Fe 10.88 (11.08) FT-IR (KBr, cm⁻¹), 3400 ν(H₂O), 3180 ν(³⁰N-H³²), 3090w, 2945, 2875w ν(C–H), 1611 ν(²³C = N²⁴), 1598 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1468 ν(¹¹N = N¹²), 1309 ν(²⁰C-O²²), 1017 ν(²⁹N-N³¹) 589 (²²O-Fe), 488 ν(²⁴N→Fe).

2.3.3. Ru³⁺-complex (4)

Color: brown, yield (77.8%), mp > 300 ^oC, (Λ_m) = 14.5 ohm⁻¹cm²mol⁻¹. Anal. for [Ru(HL)Cl₂(H₂O)₂], C₁₆H₁₇RuN₆O₃Cl₂, (513.32 g/mol): Found (calcd.) %C 37.40(37.44), %H 3.09(3.34), %N 16.43(16.37), %Cl 13.77 (13.56), %Ru 19.34 (19.69). FT-IR (KBr, cm⁻¹), 3432 ν(H₂O), 3309 ν(³⁰N-H³²), 3095w, 2995, 2870w ν(C–H), 1640 ν(²⁶C = N²⁷), 1601 ν ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1470 ν(¹¹N = N¹²), 1306 ν(²⁰C-O²²), 1012 ν(²⁹N-N³¹) 600 (²²O-Ru), 481 ν(²⁴N→Ru).

2.3.4. Co²⁺-complex (5)

Color: reddish brown, yield (67.6%), mp = 281 ^oC, Λ_m= 12.9 ohm⁻¹cm²mol⁻¹. Anal. for [Co(HL)(O)COCH₃(H₂O)₃], C₁₈H₂₂CoN₆O₆, (477.34 g/mol): Found(calcd.) %C 45.00(45.29), %H 5.03(4.65), %N 17.46(17.61), %Co 12.09(12.35). FT-IR (KBr, cm⁻¹), 3407 ν(H₂O), 3200 ν(³⁰N-H³²), 3060w, 2923, 2850w ν(C–H), 1616 ν(²³C = N²⁴), 1600 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1476 ν(¹¹N = N¹²), 1292 ν(²⁰C-O²²), 1009 ν(²⁹N-N³¹) 531 (²²O-Co), 469 ν(²⁴N→Co), 1558/1352(206) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ).

2.3.5. Mn²⁺-complex (6)

Color: brown, yield (66.7%), mp > 300 ^oC, Λ_m= 4.4 ohm⁻¹cm²mol⁻¹. Anal. for [Mn(HL)(O)COCH₃(H₂O)], C₁₈H₁₈MnN₆O₄, (437.32 g/mol): Found(calcd.) %C 49.13(49.44), %H 3.87(4.15), %N 18.89(19.22), %Mn 12.61(12.56). FT-IR (KBr, cm⁻¹), 3420 ν(H₂O), 3150 ν(³⁰N-H³²), 3060w, 2921, 2850w ν(C–H), 1613 ν(²⁶C = N²⁷), 1600 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1480 ν(¹¹N = N¹²), 1290 ν(²⁰C-O²²), 1014 ν(²⁹N-N³¹) 505 (²²O-Mn), 491 ν(²⁴N→Mn), 1542/1339(203) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ).

2.3.6. Ni²⁺-complex (7)

Color: light brown, yield (71.1%), mp = 276 ^oC, Λ_m= 14.9 ohm⁻¹cm²mol⁻¹. Anal. for [Ni(HL)(O)COCH₃(H₂O)], C₁₈H₁₈NiN₆O₄, (441.07 g/mol): Found(calcd.) %C 48.86(49.02), %H 4.02(4.11), %N 18.85(19.05), %Ni 13.06(13.31). FT-IR (KBr, cm⁻¹), 3441 ν(H₂O), 3209 ν(³⁰N-H³²), 3027w, 2920, 2861w ν(C–H), 1635 ν(²³C = N²⁴), 1607 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1465 ν(¹¹N = N¹²), 1303 ν(²⁰C-O²²), 1015 ν(²⁹N-N³¹) 500 (²²O-Co), 482 ν(²⁴N→Co), 1587/1384(203) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ).

2.3.7. Cu²⁺-complex (8)

Color: dark green, yield (76.1%), mp = 270 ^oC, Λ_m= 7.1 ohm⁻¹cm²mol⁻¹. Anal. for [Cu(HL)(O)COCH₃(H₂O)], C₁₈H₁₈CuN₆O₄, (455.90 g/mol): Found(calcd.) %C 48.95(48.48), %H 3.89(4.07), %N 14.16(14.35), %Cu 13.99(14.25). FT-IR (KBr, cm⁻¹), 3420 ν(H₂O), 3213 ν(³⁰N-H³²), 3029w, 2920, 2858w ν(C–H), 1612 ν(²³C = N²⁴), 1600 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1482 ν(¹¹N = N¹²), 1288 ν(²⁰C-O²²), 1014 ν(²⁹N-N³¹) 541 (²²O-Cu), 481 ν(²⁴N→Cu), 1569/1351(218) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ).

2.3.8. Cu²⁺-complex (9)

Color: dark green, yield (66.5%), mp = 290 ^oC, Λ_m= 11.2 ohm⁻¹cm²mol⁻¹. Anal. for [Cu(HL)Cl(H₂O)], C₁₆H₁₅CuN₆O₂Cl, (422.3 g/mol): Found(calcd.) %C 45.23(45.50), %H 3.35(3.58), %N 19.58(19.90), %Cu 14.95(15.05). FT-IR (KBr, cm⁻¹), 3413 ν(H₂O), 3162 ν(³⁰N-H³²), 3030w, 2921, 2862w ν(C–H), 1627 ν(²⁶C = N²⁷), 1607 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1477 ν(¹¹N = N¹²), 1304 ν(²⁰C-O²²), 1013 ν(²⁹N-N³¹) 511 (²²O-Cu), 456 ν(²⁴N→Cu).

2.3.9. Cu²⁺-complex (10)

Color: dark olive, yield (73%), mp = 300 ^oC, Λ_m= 21.1 ohm⁻¹cm²mol⁻¹. Anal. for [Cu(HL)NO₃(H₂O)], C₁₆H₁₅CuN₇O₅, (448.9 g/mol): Found(calcd.) %C 42.50(42.81), %H 3.27(3.37), %N 21.55(21.84), %Cu 14.01(14.16). FT-IR (KBr, cm⁻¹), 3419 ν(H₂O), 3244 ν(³⁰N-H³²), 3044w, 2931, 2854w ν(C–H), 1617 ν(²⁶C = N²⁷), 1601 ν(²⁹C = N³²)&ν(³⁰C = N³⁴), 1478 ν(¹¹N = N¹²), 1283 ν(²⁰C-O²²), 1021 ν(²⁹N-N³¹) 551 (²²O-Cu), 491 ν(²⁴N→Cu), 1408, 1363, 966 ν (NO₃)(Δ = 45).

2.3.10. Cu²⁺-complex (11)

Color: dark olive, yield (61.8%), mp=280 ^oC, Λ_m= 27.5 ohm⁻¹cm²mol⁻¹. Anal. for [Cu(H₂L)SO₄(H₂O)] C₁₆H₁₆CuSN₆O₆, (483.95 g/mol): Found(calcd.) %C 39.46(39.71), %H 3.01(3.33), %N 17.21(17.37), %S 6.18(6.62), %Cu 12.89(13.13). FT-IR (KBr, cm⁻¹), 3407 ν(H₂O), 3180 ν(³⁰N-H³²), 3027w, 2940, 2875w ν(C–H), 1629 ν(²³C = N²⁴), 1609 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1468 ν(¹¹N = N¹²), 1324 ν(²⁰C-O²²), 1011 ν(²⁹N-N³¹) 533 (²²O-Cu), 461 ν(²⁴N→Cu), 1154, 1053 ν(SO₄).

2.3.11. Zn²⁺-complex (12)

Color: light brown, yield (61.7%), mp = 284 ^oC, Λ_m= 5.4 ohm⁻¹cm²mol⁻¹, Anal. for [Zn(HL)(O)COCH₃(H₂O)₃].4H₂O C₁₈H₃₀ZnN₆O₁₀, (555.85 g/mol): Found(calcd.) %C 38.71(38.90), %H 5.24(5.44), %N 15.15(15.12), %Zn 11.55(11.76). FT-IR (KBr, cm⁻¹), 3410 ν(H₂O), 3225 ν(³⁰N-H³²), 3060w, 2937, 2855w ν(C–H), 1610 ν(²³C = N²⁴), 1600 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1465 ν(¹¹N = N¹²), 1319 ν(²⁰C-O²²), 1016 ν(²⁹N-N³¹) 546 (²²O-Zn), 474 ν(²⁴N→Zn), 1575/1388(188) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ). ¹H-NMR (DMSO-d₆, 400 MHz,): δ(ppm) = 8.62 (s, 1H, NH³⁰), 7.88 (d, 2H, C-H^8&9), 8.12 (m, 1H, CH^7&10), 7.55 (d, 1H, CH²¹), 8.71 (d, 1H, CH¹⁷), 8.92 (s, H, CH¹⁹), 9.00 (s, H, CH²⁵), 7.86 (s, H, CH²⁸), 2.57(s, 3H, CH₃), 1.98 (s, 3H, CH₃COO).

2.3.12. UO₂²⁺-complex (13)

Color: light brown, yield (62.8%), mp = 281 ^oC, Λ_m= 7.1 ohm⁻¹cm²mol⁻¹, Anal. for [UO₂(HL)(O)COCH₃(H₂O)₃].6H₂O, C₁₈H₃₄UN₆O₁₄, (796.53 g/mol): Found(calcd.) %C 27.04(27.14), %H 4.16(4.30), %N 10.43(10.55), %U 29.74(29.88). FT-IR (KBr, cm⁻¹), 3400 ν(H₂O), 3131 ν(³⁰N-H³²), 3077w, 2956, 2866w ν(C–H), 1622 ν(²³C = N²⁴), 1603 ν(²⁶C = N²⁹)&ν(²⁷C = N³¹), 1473 ν(¹¹N = N¹²), 1309 ν(²⁰C-O²²), 1017 ν(²⁹N-N³¹) 589 (²²O-U), 499 ν(²⁴N→U), 963 (O = U = O), 1568/1366(202) ν_s(CH₃COO)/ν_as(CH₃COO)(Δ).

2.4. Theoretical calculations

Based on Density Functional Theory (DFT), quantum chemical parameters was determined by Gaussian 09 Rev. A.02-MSP program [51]. The software Gauss View Rev. 5.0.9 was employed to visualize the resulted data [52]. The optimization of dihydrazone-oxime’s molecular structure in ground state was done based on DFT with Beckee3eLeeeYangeParr (B3LYP) functional levels with basis set (6-311G+(d,p)) [53,54].

2.5. Microbicidal evaluation

The assessment of in-vitro antifungal and antibacterial activities of the produced compounds (1–13) was performed utilizing agar well diffusion technique versus different fungal and bacterial species as Escherichia coli, (E. coli), Aspergillus niger (A. niger), and Bacillus subtilis (B. subtilis) [55,56]. The fungus and bacterium species were cultivated at pH 7.4 on Czapek Dox's agar (CDA), and Mueller-Hinton agar correspondingly. For A. niger, the agar plates were incubated at 28°C for 4 days, but for bacteria at 37°C for 24 h. A positive control amphotericin B (AmB) for fungi and Tetracycline (Tc) for bacteria, as well as negative (DMSO; 2% v/v) were also involved for comparison. The presence of inhibitory areas reflects the tested drugs’ antimicrobial activity. However, the three independent replicates’ average was determined and the activity index was calculated [25].

3. Results and discussion

The ligand, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-4-(o-tolyldiazenyl)phenol (H₂L, 1) was reacted with the salts of the following metallic cations VO²⁺, Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺, Zn²⁺, UO₂²⁺ leading to the construction of complexes with the subsequent general formulas [M(HL)YXZ(H₂O)].nH₂O (where M = Fe^3 +(3), Ru^3 +(4), Co^2 +(6), Mn^2 +(7), Ni^2 +(5), Cu^2 +(8–10), Zn^2 +(12) and UO₂^2 +(13), X = Cl, CH₃COO or NO₃ and Y = Cl, or H₂O, Z = H₂O and n = 0, 2 or 6), [Cu(H₂L)SO₄(H₂O)] and [VO(H₂L)SO₄(H₂O)₂]. The prepared compounds’ structure was elucidated analytically, theoretically, thermally and by several spectroscopical tools, such UV-Vis., NMR, and FT-IR as well as mass spectrometry. The data of different measurements and theoretical calculations were recorded in the experimental section and Tables 1–4 as well as Tables S1–S4. The postulated molecular formulae were supported by elemental analysis and mass spectrometry, which also revealed that all complexes formed in a (1L:1M) molar ratio (Figures 2–3). All complexes have molar conductance value ranged from of 4.1–34.9 Ω⁻¹mol⁻¹cm², indicative to the non-electrolyte nature of the synthetics (2–13) and the Cl⁻, NO₃⁻_, SO₄²⁻ or CH₃COO^- are directly connected to the metallic cations [57]. A few complexes’ noticeably high values may be caused by the partial solvolysis by DMSO molecules [57].

Figure 2. Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺, Zn²⁺, and UO₂²⁺ complexes’ Structure.

Figure 3. VO²⁺, and Cu²⁺ complexes’ structure.

Table 1. Data of prepared compounds’ EAS.

Compound	Bands (nm) in DMSO	Electronic transition	µeff (BM)	Geometry
Azo-Schiff base	260, 283, 313, 333	π→π*, n→π*	^–
VO^2+complex (2)	262, 285, 327, 365, 430, 470, 809, 1072	^2B2g(dxy)→^2Eg(dxz,dyz)(ν1), ^2B2g(dxy)→^2B1g(dx2-y2)(ν2) ^2B2g(dxy)→^2A1g(dz2)(ν3)	1.67	Distorted octahedral
Fe^3+complex (3)	263, 287, 310, 341, 410, 445,585, 809	(ν3)^6A1g(S)→^4A1g(G) ^4E1g(G) (ν2)^6A1g(S)→^4T2g(G) (ν1)^6A1g(S)→^4T1g(G)	5.88
Ru^3+complex (4)	267, 281, 369, 407, 425, 539	(Lπy→T2g) ^2T2g→^2A2g	1.66
Co^2+complex (5)	266, 307, 391, 436, 568, 1090	(v2)^4T1g(F)→^4A2g (v1)^4T1g(F)→^4T2g(F)	4.46
Mn^2+ complex (6)	265, 310, 345, 394, 434, 491, 545	π→π*, n→π*	5.78	tetrahedral
Ni^2+complex (7)	265, 310, 405, 440, 663, 940	(ν3)^3T1g(F)→^3T1g(P) (ν2) ^3T1g(F)→^3A2g(P)	3.08
Cu^2+complex (8)	265, 321, 366, 428, 629, 1093	(υ3)^2B1g→^2Eg(dx2-y2→dxy) (υ2)^2B1g→^2A1g (dx2-y2→dz2) (υ1)^2B1g→^2B2g (dx2-y2→dyzdxz)	2.2	Square planar
Cu^2+complex (9)	266, 309, 371, 454, 650, 1089		1.88
Cu^2+complex (10)	269, 295, 358, 485, 622, 1083		1.97
Cu^2+ complex (11)	265, 310, 405, 463, 670, 1086		2.11
Zn^2+ complex (12)	264, 290, 341, 425	π→π*, n→π*	Dia.	–
UO2^2+ complex (13)	265, 352, 373, 423, 486	LMCT O→U	Dia.	–

Table 2. ESR parameters of Cu²⁺complexes.

Complex No.	(8)	(9)	(10)	(11)
g||	2.253	2.214	2.230	2.260
g┴	2.60	2.42	2.054	2.048
giso	2.124	2.099	2.113	2.119
A|| × 10^−4 (cm^−1)	179	186	167	201
A┴×10^−4 (cm^−1)	38.2	27.9	28.3	19.1
Aiso ×10^−4 (cm^−1)	82.2	84.6	71.6	75.9
g||/A|| (cm^−1)	126	119	133.8	112.7
G	4.19	5.3	4.41	5.62
ΔExy(cm^−1)	15898	15384	16077	14925
ΔExz (cm^−1)	23364	22026	20619	21598
K||	0.776	0.701	0.746	0.762
K⊥	0.905	0.729	0.804	0.774
K	0.864	0.720	0.785	0.77
α^2	0.786	0.754	0.665	0.761
β^2	0.70	0.48	0.27	0.68
γ	0.75	0.63	0.75	0.66

Table 3. The prepared compounds’ various quantum chemical parameters.

No.	EHOMO (a.u.)	IP (a.u.)	ELUMO (a.u.)	EA (a.u.)	ΔE (a.u.)	χ(a.u.)	η (a.u.)	σ (a.u.^−1)	Pi = −χ (a.u.)	ω (a.u.)	ΔNmax
H2L	−0.19193	0.1919	−0.09025	0.0903	0.1017	0.1411	0.0508	19.67	−0.1411	0.1958	−2.78
(2)	−0.2112	0.2112	−0.1329	0.1329	0.0784	0.1720	0.0392	25.52	−0.1720	0.3776	−4.39
(3)	−0.1987	0.1987	−0.1197	0.1197	0.0791	0.1592	0.0395	25.29	−0.1592	0.3205	−4.03
(4)	−0.1898	0.1898	−0.1008	0.1008	0.0890	0.1453	0.0445	22.48	−0.1453	0.2374	−3.27
(5)	−0.20372	0.2372	−0.14446	0.1445	0.0593	0.1741	0.0296	33.73	−0.1741	0.5114	−5.86
(6)	−0.1665	0.1665	−0.1070	0.1070	0.0595	0.1368	0.0297	33.63	−0.1368	0.3145	−4.60
(7)	−0.1442	0.1442	−0.0848	0.0848	0.0594	0.1145	0.0297	33.68	−0.1145	0.2206	−3.68
(8)	−0.2161	0.2161	−0.1158	0.1158	0.1003	0.1659	0.0501	19.95	−0.1659	0.2746	−3.31
(9)	−0.2172	0.2172	−0.1178	0.1178	0.0994	0.1675	0.0497	20.12	−0.1675	0.2822	−3.37
(10)	−0.2175	0.2175	−0.1174	0.1174	0.1000	0.1674	0.0500	19.99	−0.1674	0.2802	−3.35
(11)	−0.2112	0.2112	−0.14346	0.1435	0.0557	0.1723	0.0289	17.32	−0.1723	0.5144	−5.97
(12)	−0.2029	0.2029	−0.1222	0.1222	0.0806	0.1626	0.0403	24.80	−0.1626	0.3277	−4.03

Table 4. Prepared compounds’ energetic properties.

compounds	EHOMO (a.u.)	ELUMO (a.u.)	Total Energy (a.u.)	binding Energy (a.u.)	Dipole Moment (Debye)
H2L	−0.19193	−0.09025	−0.1017	−1021.4	3.29
VO^2+complex (2)	−0.2112	−0.1329	−0.784	−2893.1	20.9
Fe^3+complex (3)	−0.1987	−0.1196	−0.0791	−3358.4	6.34
Ru^3+complex (4)	−0.1898	−0.1008	−0.0890	−6506.6	6.53
Co^2+complex (5)	−0.20372	−0.14446	−0.0593	−2861.8	7.55
Mn^2+complex (7)	−0.1665	−0.10702	−0.0594	−2477.3	7.40
Ni^2+complex (6)	−0.1442	−0.0848	−0.0594	−2708.3	3.68
Cu^2+complex (8)	−0.21633	−0.11469	−0.1016	−2996.8	6.20
Cu^2+complex (9)	−0.21721	−0.11780	−0.0994	−3198.5	5.56
Cu^2+complex (10)	−0.21746	−0.11742	−0.1000	−3018.6	6.25
Cu^2+complex (11)	−0.20119	−0.14346	−0.0557	−3438.0	15.4
Zn^2+complex (12)	−0.20288	−0.12224	−0.0806	−3258.2	9.42

3.1. Spectroscopical studies

3.1.1. FT-IR spectra

In the experimental section, the FT-IR spectral information (experimental and calculated) (Figures S1–S4) was presented and clarified as follow: the literature values of OH and NH stretching vibrations are found to be in 3400–3600 and 3150–3300 cm⁻¹ ranges [58]. In the studied Azo-Schiff base molecule (H₂L, 1), there is a hydroxyl υ(O-H) and υ(N-H) bond. The stretching vibration bands of these bonds have been calculated at 3750, 3275 and 3230 cm⁻¹. In the experimental spectrum there are two bands at 3484, 3233 cm⁻¹ in the ligand’s FT-IR spectrum may be allocated to the hydroxyl υ(O-H) and υ(N-H) groups successively. The molecule of Azo-Schiff base (H₂L, 1), includes three azomethine bonds, the literature values of C = N stretching vibrations are between 1500 and 1700 cm⁻¹ [59]. In the studied Azo-Schiff base (H₂L, 1), stretching vibration bands have been calculated at 1690, 1640 and 1600 cm⁻¹. In the experimental spectrum, vibrations observed at 1660 and 1605 cm⁻¹ have been assigned to the C = N stretching bands of Schiff base and triazole moieties. In the experimental spectrum, the medium bands which observed at 1475, 1267 and 1000 may be ascribed to the of azo (N = N), phenolic (C–O) linkages and (N-N) linkage of a triazole moiety successively with theoretical values at 1490, 1100 and 1255 cm⁻¹. The coordination mode of azo-Schiff base (H₂L, 1) can be ascertained by comparing the complexes’ FT-IR to those of un-chelated azo-Schiff base (H₂L, 1). From this comparison we can conclude that the azo-Schiff base (H₂L,1) interacted with different metallic cations in one of the following two ways. Firstly, it interacted with Fe³⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, Cu²⁺(8–10), Zn²⁺ and UO₂²⁺ as uni-negative bidentate ligand linked covalency the metallic cations via deprotonated phenolic hydroxyl group (O²²) and chelated via the azomethine nitrogen atoms (N²⁴). The following observations served as confirmation of this chelation method; (i) The disappearing of the phenolic hydroxyl group’s stretching frequency with simultaneous positive shift of phenolic linkage (C–O) by 16–52 cm⁻¹ [60]. (ii) the negative shift in the Schiff base azomethine group’s stretching frequency (²³C=N²⁴) by 20–50 cm⁻¹. (iii) the appearance of new stretching frequencies in the ranges of 500–600 and 456–491 cm⁻¹ may be inductive to the ν(²²O-M) and ν(²⁴N→M) successively [61]. In the second manner, the azo-Schiff base (H₂L,1) interacted with VO²⁺ and Cu^2 +(11) ions as a neutral bidentate ligand in which the coordination took place via the protonated hydroxyl (²²OH) and Schiff base azomethine’s (²⁴N) groups. The following observations served as confirmation of this chelation method; (i) the simultaneous shift of the phenolic linkage (C-OH) and the phenolic hydroxyl (OH) group's stretching frequencies [60]. (ii) The negative shift in the Schiff base azomethine group’s stretching frequency (²³C = N²⁴) by 27 and 31, respectively cm⁻¹. (iii) The emergence of new stretching frequencies at 577, 533 and 491, 461 cm⁻¹ which could be inductive to the ν(²²O→M) and ν(²⁴N→M) successively [61]. In Co²⁺, Mn²⁺, Ni²⁺, Cu^2 +(8), Zn²⁺ and UO₂²⁺ complexes, the appearance of two distinctive stretching frequency bands in the of 1542–1587 cm⁻¹ and 1339-1388 cm⁻¹ ranges may be ascribed to υ_as.(COO^-) and υ_sy.(COO^-) successively, signifying that the acetate group participated in the chelation with metallic cations. Based on the difference’s magnitude between the υ_as(COO^-) and υ_sy.(COO^-), the coordination mode of acetate group can be determined. The difference between the ν_as and ν_s stretching frequency bands (Δ =  ν_as-ν_s) which varied from 188 to 218 cm⁻¹, indicating that the acetate anion linked with metallic cations in a monodentate manner [62,63]. While the two bands at 1155, 1061 and 1154, 1053 cm^−1, in the sulfato complexes’ FT-IR spectra successively imputing to the mono-chelated sulphate group which was corroborated by their low values of molar conductivity [2,64]. The nitrato Cu^2 +(10) complex displayed three bands at ν₅(1408), ν₁(1363) and ν₂(966) implying that the nitrate anion participated in the coordination process. The difference between ν₅ and ν₁ stretching frequency bands (ν₅-ν₁) is 45 cm⁻¹ indicating that the NO₃⁻ anion is linked with the Cu²⁺ cation as a uni-dentate chelator [62,65]. The new stretching frequency band at 963 and 970 cm−1 in the VO²⁺ and UO₂²⁺ complexes FT-IR’ spectra might be ascribed to the ν(O = U = O) and ν(V = O) successively [61,66]. According to literature, the methyl groups’ C-H vibrations appear around 2900 cm1, while the aromatic’ C-H vibrations are commonly seen at about 3000 cm1. So, the feeble bands observed in the 3095–3027 cm1 range may be ascribed to the symmetric vibrational bands of aromatic protons, while the bands observed in the 2920–2995 cm1 and from 2854 to 2875 cm1 regions may be ascribed to pure asymmetric and symmetric methyl protons’ vibrational bands [67].

3.1.2. Nuclear magnetic resonance (NMR)

To get essential info around the hydrogen bonding and coordination mode of the azo-Schiff base (H₂L,1), as well as to identify its structure. The ¹³C- and ¹H-NMR spectra were obtained in DMSO-d₆ and summarized in section 2.2. In ¹HNMR the disappearance of aldehydic proton of the reactant at ≈ 9.9 ppm indicates to the formation of imine group between the triazole’s amine and the aldehydic moiety of the diazenyl reactant. In addition, the disappearance of aldehydic carbon at ≈190 ppm of the reactant in ¹³CNMR and it is replaced with up field carbon at 166.31 ensured the formation of imine moiety in the azo-Schiff base [68,69]. In ¹HNMR (Figure S5), the compound has three exchangeable protons, two protons at 8.35, 14.37 which are assigned to the protons of NH implying to the anti- and syn-configurations of the azo-Schiff base around. Therefore, the ligand formed is a mixture of two configurations as shown in Figure 4. From the size of the two signals, we can conclude that the most abundant one is the anti-structure [70]. The third signal appeared at 12.619 assigned to phenolic hydroxyl proton. This result was supported by the fact that their intensity decreased when the solvent was changed from DMSO to DMSO + D₂O, which coincided with the downfield value of C-20 at 166.31 ppm in the benzene ring because of the existence of OH at C-20. ¹HNMR also indicated the presence of one downfield proton at 9.56, which could be ascribed to the methine proton produced from the imine formation. Additionally, the CH signals overlapped in the 7.31–9.13 ppm range imputing to the two phenyl rings’ aromatic protons. The noted chemical shift of H¹⁷, H^8&9 (δ = 8.819, 8.395 ppm) is a characteristic downfield value of protons in o-position for the diazinyl moiety, whereas the downfield shift of the singlet peak of H¹⁹(9.13) and the up-filed shift of H²¹ (7.31) suggested the presence of the hydroxyl group at ortho and meta positions to these two protons, respectively [71]. The proton of the triazole ring appears at 8.173 ppm. The Zn²⁺ complex’s ¹H-NMR spectrum revealed that the chemical shift of the phenolic hydroxy proton (³³H-O²²) is disappeared indicating that the hydroxyl group participated in the metallic ion’s chelation in its deprotonated form which endorsed the proposed chelation mode. The azo-Schiff base’s ¹³C-NMR spectrum is illustrated in Figure S6. It showed two chemical shifts 159.58 and 146.6 ppm which could be assigned to two carbons of triazole ring (²⁶C&²⁷C) while the chemical shift of C²⁰ (δ = 166.31 ppm) indicates the presence of a hydroxyl group attached to this carbon. The chemical shifts of carbons C², C⁴, C¹⁴, C¹⁵ (δ = 122.47, 125.51, 125.51, 128.29 ppm) are typical values for carbon attached to C¹³ (δ = 144.12 ppm) and C³ (δ = 150.63 ppm) connected to azo group (N = N). However, the peaks observed at 134.16, 140.47, 113.79, and 117.21 ppm were allocated to the aromatic carbons C^1&5, C⁶, C¹⁶, and C¹⁸, respectively. The peak appearing at 19.57 ppm could be due to the methyl carbon attached to the benzene ring. The assignment of carbon and protons is confirmed by comparing with similar compounds [72].

Figure 4. The Anti- and Syn-structures of azo-azomethine ligand (H₂L, 1).

3.1.3. Azo-Schiff base’s mass spectrometry (MS)

The azo-Schiff base’s MS (Figure S7) proved the expected composition of it with molecular ion peak (M⁺) at 306 m/z in EI-mass. Additionally, many fragments appeared in the mass spectrometry conforming the structure. The ligand revealed several fragments at 292, 290, 276, 239, 216, 212, 188, 119, 96, 92, 77 and 69 m/z. Which could be clarified as follows. The two fragments (A&B) at m/z 292 and 290, may be because of the missing of hydroxyl or methyl group from the azo-Schiff base molecule respectively which consequently lose methyl or hydroxy group resulting to the fragment (C) formation (m/z = =276). While the 2-(iminomethyl)-4-(o-tolyldiazenyl)phenol fragment that resulted from the loss of a triazole moiety may be attributed to the d-fragment (m/z = 239) which lose C = NH moiety resulting to the fragment (E) formation (m/z = 212) that directly lose a phenolic moiety forming 1-(o-tolyl)-2λ²-diazene (F). This fragment consequently loses azo-linkage (N = N) forming toluene ending with the formation of phenyl moiety. The fragments A&B may lose a phenyl ring resulting in the formation of G-fragment (z = 216) which directly loses a N₂ forming a fragment (H, m/z = 188). The fragment (H) lost a phenolic moiety forming a fragment (I, m/z = 96) which directly lost a C-NH moiety forming a triazole moiety (Figure S8). The MS of [Cu(HL)Cl(H₂O)] (Figure S9) revealed a weak molecular ion peak at 422 with low intensity and another peak at 405 with low intensity corresponding exactly to [Cu(HL)Cl]. Also, there are peaks at 369 and 306 which may correspond to [Cu(HL)] and (HL). On the other hand, the MS of [Zn(HL)((O)COCH₃)(H₂O)₃].4H₂O (Figure S10) showed a peak at 483 (7%) corresponding exactly [Zn(HL)((O)COCH₃)(H₂O)₃] with no peaks higher than this one. This species showed peaks at 460(2%), 448(3) and 430(3) which formed after the successive removal of coordinated water molecules. The medium peak observed at 370 (15%) may correspond to [Zn(HL)] while the peak at 306 may be due to ligand molecule after the elimination of Zn atom.

3.1.4. Magnetic moment (µ_eff) and electronic absorption spectroscopic (EAS) measurements

The azo-Schiff base’s EAS and that of metallic complexes (1–13) were determined in DMSO and itemized in Table 1 and shown in Figures S11–S14. Ligand’s EAS revealed two sets of bands in the UV and visible regions. The bands that appeared at the shortest wavelengths (283 and 260 nm) could be attributed to the π→π* and π→π* transitions, which were attributed to intra-ligand transitions and transitions in the phenyl moieties [73]. While the two bands at 313 and 333 could be inductive to the n→π* transition of Schiff base and triazole azomethine groups [73]. The complexes’ n-π* transitions were shifted somewhat compared to the free ligand, which is most likely due to the chelation of azomethine-nitrogen with the metallic ions [74]. In hexa-bonded VO²⁺ complex the ground state ²T_2g and ²E_g splits into ²B_2g, ²E_g and ²B_1g, ²A_1g, respectively thus in octahedral field there are three spin allowed transitions corresponding to ²B_2g(d_xy)→²E_g(d_xz,d_yz)(ν₁), ²B_2g(d_xy)→²B_1g(d_x2-y2)(ν₂) and ²B_2g(d_xy)→²A_1g(d_z2)(ν₃) successively [75]. So the peaks that appeared at 1048, 809, and 470 nm in the EAS of the VO²⁺ complex (2) are imputable to the ²B_2g(d_xy)→²E_g(d_xz,d_yz)(ν₁), 2B2g(dxy)→2B1g(dx2-y2)(ν2), and ²B_2g(d_xy)→²A_1g(d_z2)(ν₃) transitions successively [76]. These transitions are consistent with a distorted octahedral structure for the VO²⁺ cation (Figure 3) [77,78]. This VO²⁺-complex has an eff value of 1.67 BM, which is consistent with the spin-only value [77,78]. In the hexa-coordinated Fe³⁺ complex, the ground state (⁶S) splits into 6A1g, 4T1g (P), ⁴T_2g, 4A1g, and ⁴E_1g. Thus, in the octahedral field, there are three spin-allowing bands relating to the following 6A1g(S)→4T1g(G) (∼950 nm), 6A1g(S)→4T2g(G) (∼600 nm), and ⁶A_1g(S)→⁴A_1g(G) ^{4E1g(G) (∼450 nm)} transitions [79]. According to this explanation, the EAS for the Fe³⁺ complex (3) demonstrated three peaks at 898, 585, and 445  nm, referring to the above transitions, which are well-suited to the distorted octahedral geometry (Figure 2) [75]. The µ_eff value of Fe³⁺ complex (3) at room temperature was found to be 5.88 BM inductive to a five-unpaired electron system, which is well-suited to high-spin octahedral Fe³⁺ cations [76]. The Ru³⁺ complex exhibits a magnetic moment value of 1.66 BM, which is compatible with a spin-only value for low-spin Ru³⁺ species [65]. The low-spin Ru³⁺ complex has a ⁵t_2g arrangement, with the first excited doublet levels being 2T1g and ²A_2g in the order of decreasing energy that arise from the ⁴t_2g→¹e_g configuration. As a result, two peaks that correspond to ²T_2g→²T_1g and ²T_2g→²A_2g are feasible. The Ru³⁺ complex’s EAS (4) nevertheless demonstrated one band at 533 that may be ascribed to the ²T_2g→²A_2g transition. While the band at 425 nm may be the result of a charge transfer transition (LMCT) from the filled chelator (HOMO) orbital to the single-busied ruthenium t₂ orbital (LUMO),. The band position is similar to that of other octahedrally structured Ru³⁺ complexes [77,80]. In octahedral symmetry for Co²⁺ complexes, ground state (⁴F) splits into ⁴A_2g, ⁴T_2g and 4T1g (P). Thus, in the octahedral field, three spin-allowing transitions correspond to the (v1) 4T1g (F)→4T2g (F), (v2) 4T1g (F)→4A2g, and (v3) 4T1g (F)→4T1g (p) [77,81]. According to this description, the EAS for Co²⁺ complex (5) showed two peaks at 568 and 1090 nm imputable to the (v2)4T1g(F)→4A2g and (v₁)⁴T₁g(F)→⁴T_2g(F) transitions successively, which is well-suited with an octahedrally structured structure (Figure 2) [77,81]. The distortion of this structure is suggested based on its small value of ν₁/ν₂ (1.92) in comparison to the typical range for octahedrally Co²⁺ complexes (1.95–2.48). This is coherent with the broadness of the bands, which is attributed to the envelope of the electronic transitions from 4Eg (4T1g) to ⁴E_g and ⁴B_2g components of ⁴T_2g in a tetragonally distorted octahedral setting [82]. The Co²⁺ complex (5) resembles a high-spin Co²⁺ ion (d⁷) where its µeff is equal to 4.46 BM, which is equivalent to three unpaired electrons [83]. The Mn²⁺ complex’s EAS showcased feeble bands at 545 and 491  nm, which are attributed to ⁶A_1g→⁴T_1g(4G) (υ1) and ⁶A_1g→⁴T₂g(υ₂) transitions successively, corresponding to a tetrahedral geometry around the Mn²⁺ ion (Figure 2) [84]. The third and fourth bands ⁶A_1g→⁴E_g(G), ⁴A_1g(G) (v3), and ⁶A_1g→⁴T_2g(D)(v₄) do not appear here because of the broadening of the bands [77,85]. This Mn²⁺ complex has a µ_eff value of 5.78 BM, which is consistent with a system having five unpaired electrons and is in close proximity to a high-spin Mn²⁺ ion (d⁵)[86]. In the tetrahedral symmetry of the Ni²⁺ complex, the ³T_2g state term originating from the ³F level becomes the ground state. Thus, in the tetrahedral field, there are three spin allowable bands comparable to the next (v₃)³T_1g(F)→³T_1g(P) (625), (ν₂)³T_1g(F)→³A_2g(P) (1000), and (ν₁)³T_1g(F)→³T_2g(F) (2000nm) transitions. The Ni²⁺ complex’s EAS exhibited two peaks at 940, 663  nm, imputable to the ³T_1g(F)→³A_2g(P)(ν₂) and ³T_1g(F)→³T_1g(P)(v₃) transitions successively, which is well-suited with a tetrahedral geometry (Figure 2) [77,87]. The 3rd band, which is attributed to the (ν₁)³T_1g(F)→³T_2g(F) (2000nm), is not noticeable because it is out of spectrophotometer scale. The Ni²⁺ complex (7) was found to have an µ_eff value of 3.08 BM at room temperature, which is within the typical range for high-spin tetrahedral Ni²⁺ complexes and rules out the idea of a square planar structure [88]. The EAS of Cu²⁺ complexes (8–11) showed three bands in the ranges 1083–1093 and 622–670, 428–485 nm. These bands are ascribable to the following transitions: (υ₁)²B_1g→²B_2g (d_x2-y2→d_yzd_xz), (υ₂)²B_1g→²A_1g (dx2-y2→dz2), and (υ3)2B1g→2Eg (dx2-y2→dxy), which are typical for square planar Cu²⁺ complexes (Figures 2–3) [77]. The Cu²⁺ complexes’ magnetic moment values range from 1.88 to 2.20 BM, which are consistent with the existence of one unpaired electron of Cu²⁺ ions in a square planar structure. Two bands were visible on the EAS of the UO₂²⁺ complex at 486 and 423 nm. Charge transfer from the UO₂²⁺ moiety's oxygen to the U⁶⁺ ion's f-orbital (O→U) is believed to be the cause of the lower energy band (486 nm). This band boarding suggests that the energy of the two O-U charge transfer transitions for oxo cations are different [89]. However, the higher energy band is thought to be responsible for the charge transfer from the ligand to the UO₂²⁺ ion (LMCT)[90].

3.1.5. Cu²⁺ complexes ESR spectrum

The data of Cu²⁺-complexes ESR spectra was tabulated in Table 2. It showed anisotropic signals with g_|| values ranged from 2.214 to 2.26, g_⊥ values ranged from 2.042 to 2.060 and g_iso values ranged from 2.099 to 2.124 confirming to a d⁹ configuration denoting to axially symmetric anisotropy type of a d_(x2-y2) ground state [91,92]. The g-values for Cu²⁺-complexes (7–8) and (10–11) supported that they have one of the following geometries an octahedral or square planar while the trend g_||>g_┴>2.0023 suggesting that these geometries could be deformed [93]. The g-value can be correlated to the parameter term of exchange interaction G by equation, G = (g_||-2)/(g_┴-2) [92]. According to Hathaway if G is less than 4.0, a considerable exchange coupling is present but if G is more than 4.0, exchange coupling is negligible. The Cu²⁺ complexes G-values equal to 4.19, 5.30, 4.41 and 5.62, respectively implying that the exchange interaction between Cu²⁺ ions is negligible [94]. Also, the g_||/A_|| values can be serve as a diagnostic of stereochemistry [95]. According to literature, for square-planar complexes the g_||/A_|| values ranged from 105 to135 cm⁻¹ but for tetragonally deformed complexes ranged from 150 to 250 cm⁻¹. The g_||/A_|| value for the Cu²⁺ complexes equal to 12.97, 118.97, 133.84 and 112.71, respectively. These data confirmed that the Cu²⁺ complexes have a square planar. If the g_||< 2.3 the environment is mainly covalent while if g_||> 2.3 the environment is essentially ionic [96]. The g_||-values for the Cu²⁺ complexes under study were ranged from to 2.214 to 2.26, implying to the covalency nature of bonds in these complexes [96]. The relationship between the g-values and the orbital reduction factor (k) can be established through the following equations [97,98]. $\begin{array}{rl}{{\text{K}}_{||}}^2& =\left({\text{g}}_{||}-2.0023\right)\mathrm{\Delta }{\text{E}}_{\text{xy}}/8{\lambda }_o\\ {{\text{K}}_{\mathrm{\perp }}}^2& =\left({\text{g}}_{\mathrm{\perp }}-2.0023\right)\mathrm{\Delta }{\text{E}}_{\text{xz}}/2{\lambda }_o\\ {\text{K}}^2& =\left({{\text{K}}_{||}}^2+2{{\text{K}}_{\mathrm{\perp }}}^2\right)/3\end{array}$where ΔE_xz and ΔE_xy are the electronic transitions ²B_1g→ ²B_2g and ²B_1g→ ²E_g, sequentially, λ_o = −828 cm⁻¹ is the spin orbit coupling of free Cu²⁺ ion. If k = 1 the environment is ionic, but if k < 1 environment is covalent. The calculating values of k is less than unity (0.72–0.86) suggesting to the covalency nature of these complexes which is in agreeable with the results attained from the g_|| values [98]. The value of in-plane σ-bonding parameters α_Cu² can be appreciated according to Kivelson and Neiman [96]. ${\alpha }^2=\left(g_{\parallel }-2.0023\right)+\frac{3}{7}\left(g_{\mathrm{\perp }}-2.0023\right)-\left(\frac{A_{\parallel }}{P}\right)+0.04$where P is the free ion dipolar term equals 0.036, $A_{\parallel }$ is the parallel coupling constant expressed in cm⁻¹. If the value of α_Cu² = 1, the bond could be pure ionic but if α_Cu² = 0.5, the bond could be pure covalent [91,93]. The values of α_Cu² for Cu²⁺-complexes were found to be range from 0.665 to 0.787 suggesting to covalency nature of in-plane σ-bonding [99]. The out-of plane and in-plane π-bonding coefficients γ² and β² can be estimated by published equations [97,100]. $\begin{array}{rl}{\alpha }^2{\gamma }^2& ={{\text{K}}_{\mathrm{\perp }}}^{\mathbf{2}}\\ {\alpha }^2{\text{β}}^2& ={{\text{K}}_{||}}^2\end{array}$The Cu²⁺ complexes γ² and β² values were found to be ranged from 0.27 to 0.70 and from 0.63 to 0.75 successively, demonstrating that the in-plane and out-of-plane π-bonding have a covalency nature [63,97].

3.1.6. Thermogravimetric analysis

By measuring the weight loss as a function of temperature, the TGA analysis can be used to not only determine the thermal stability of complexes under investigation but also to give essential information about the water molecules’ nature in the complexes’ skeleton. As record in Table S1 and shown in Figures S15–S17, the isolated complexes’ thermograms recorded in temperature ranged from room temperature to 800°C. The complexes’ thermogram showed noteworthy continuous phases of decomposition in the temperature range of 41–632°C, which culminated in the creation of residues made of metal oxide. The subsequent steps of decomposition may be explained as follows. In the thermograms of Fe³⁺, and Zn^2+, four mass loss steps were seen. The first stage of the degradation took place between 34 and 133°C, with weight losses of 6.71, and 1255% (calcd at 7.15, and 12.96%) related to the liberation of two, or four hydrated water molecules, respectively. The second decaying stage occurred in the range of temperature of 142–248°C with loss of two or three chelated water molecules, with a mass loss of 6.49, or 10.01% (calcd 7.15, 9.72%) successively. The third stage of decaying occurred in the 234–354°C temperature range which was accompanied by loss of two chloride or one acetate anion, which resulted in mass losses of 13.87, and 10.34% (calcd. 14.07, 10.69%). The fourth phase of decaying took place between 311 and 632°C, with mass losses of 52.91, and 50.89% (calcd at 53.42, 51.98), corresponding to the complete complexes’ breakdown with the formation of Fe₂O₃, and ZnO, respectively. The VO²⁺, Ru³⁺, Co²⁺, Mn²⁺, Ni²⁺, and Cu²⁺ complexes’ thermograms revealed three weight lose steps. In the temperature range 95–310°C, the first breakdown step involves the loss of one, two, or three chelated water molecules associated with weight falling of 4.05, 6.88, 11.00, 4.41, 3.81, 3.55, 4.55, 3.88 and 3.33% (calcd. 3.70, 7.02, 11.32, 4.12, 4.08, 4.04, 4.27, 4.01 and 3.72%). The second breakdown step occurred in temperature range 194–380°C associated with falling in weight equal to 19.41, 13.46, 11.87, 13.76, 13.41, 13.76, 8.41, 13.41 and 19.76% (calcd. 19.70,13.81, 12.45, 13.59, 13.48, 13.33, 8.40, 13.81 and 19.84%) referring to the losing of sulfate, chloride, acetate, or nitrate anions. The third decomposing stage happened in the temperature range 280–543°C with losing in weight equal to 58.96, 55.43, 60.89, 66.56, 63.77, 66.01, 67.76, 64.76 and 58.20% (calcd. 57.95, 54.80, 60.53, 66.07, 65.51, 64.79, 68.50, 64.56 and 60.00%) leading to the complete complexes’ breakdown leaving the corresponding metal oxide (V₂O₅, Ru₂O₃, CoO, MnO, NiO, and CuO) [101–104]. While in the TG-graph of UO₂²⁺, three mass loss steps were seen. The first stage of the degradation took place between 50 and 160°C, with weight losses of 13.81% (calcd at 13.57%) related to the liberation of six hydrated water molecules. The second decaying stage occurred in the range of temperature of 160–255°C with loss of chelated water molecules two and one acetate ion, with a mass loss of 14.50% (calcd 14.25%). The third stage of decaying occurred between 260 and 528°C, with mass losses of 35.76% (calcd 36.28%), corresponding to the complete complex’s breakdown with the formation of UO₃.

3.2. Density functional theory elucidation

3.2.1. Geometrical optimization

Figures 5 and S18–S20 depict the completely optimized structure for the azo-Schiff (H₂L, 1) and its chelated complexes. The complexes’ chosen bond length and bond angle values showed distorted octahedral, square planar or tetrahedral geometry around the metallic cations (Tables S2–S3). In complexes the deprotonated/protonated phenolic oxygen and azomethine nitrogen atoms of azo-Schiff base (H₂L) bonded the metallic cations. They displayed a slight shrinkage at bond lengths of the azomethine linkage (²³C=N²⁴) and the protonated phenolic hydroxy linkage (²⁰C−²²OH) as in VO²⁺ and Cu²⁺ (11) complexes. But they displayed an elongation at bond lengths of deprotonated phenolic hydroxy linkage (²⁰C-O²²). The bond length between the metallic cation and deprotonated phenolic oxygen atom was found to be in the 1.80–1.91 Å range. While the bond length between the protonated phenolic oxygen or azomethine nitrogen atoms and metallic cations were found to be in the 2.68496–2.85694 Å range. In hexa-coordinated VO²⁺, Fe³⁺, Ru³⁺, Co²⁺ and Zn²⁺ complexes, the two axial sites were busied by two chelated water molecules while in Fe³⁺, Ru³⁺, Co²⁺ and Zn²⁺ complexes but in VO²⁺ complex the axial positions occupied by one chelated water molecule and vanadyl oxygen. While the four equatorial locations were inhabited by phenolic oxygen, azomethine nitrogen, two chloride anions in Fe³⁺, Ru³⁺, but in complexes VO²⁺, Co²⁺ and Zn²⁺ complexes, the four equatorial locations were busied by phenolic oxygen, azomethine nitrogen and chelated molecule molecules. In tetrahedral Ni²⁺ and Mn²⁺ the four locations of tetrahedron were inhabited by azomethine nitrogen, phenolic oxygen, acetate anion and chelated water oxygen. In square planar Cu²⁺ complexes the four corners of square were occupied by azomethine nitrogen, phenolic oxygen, anion (Acetate, Chloride, nitrate, or sulphate) and chelated water oxygen. The angles of bond in the coordination sphere of complexes were indicative to distortion of the structure. The geometric index for four-coordinate complexes, τ₄ was calculated by τ₄ = 360° − (α+ β)/141°. Where α and β are the two greatest valance angles of coordination center. According to literature, for square-planar geometry the τ₄ values equal 0 but for tetrahedral geometry the τ₄ values equal 1. While intermediate geometry, involving seesaw and trigonal pyramidal, will be ranged from 0 to 1. The τ₄ values for the four-coordinate Ni²⁺ Mn²⁺ and Cu²⁺ complexes equal to 0.91, 0.93, 0.44, 0.39, 0.37 and 0.51, respectively. These data confirmed that the Ni²⁺ and Mn²⁺ complexes have predominately distorted tetrahedral geometry while Cu²⁺ complexes are predominately distorted square planar geometry [105,106].

Figure 5. The Optimized structure for azo-Schiff base (H₂L).

3.2.2. Molecular parameters

Basing on Koopman theorem, The azo-Schiff base’s quantum chemical parameters and that of its metal chelates (2–12) were calculated and their data were itemized in Tables 3–4 [107]. The azo-Schiff base’s molecular orbital and that of its chelated complexes (LUMO and HOMO) were demonstrated by DFT/B3LYP/6-311G+(d,p) calculation and it is shown in Figure 6. The negatively and positively charged areas were shown using red and green colors successively. Figure 6 demonstrated that the HOMO molecular orbital was spread nearly over the azo moiety including one phenyl ring, some carbons atoms of the second phenyl ring and phenolic oxygen whereas the LUMO molecular orbital was spread over the triazole moiety, azo linkage, one phenyl ring and some carbons of the second phenyl ring as well as phenolic oxygen and azomethine nitrogen. The energies of the HOMO and LUMO as well as the nearby molecular orbitals are negative, demonstrating the stability of the azo-Schiff base molecule [108]. The azo-Schiff base’s chelation with metallic cations happened via the phenolic oxygen atom (O²²), and azomethine nitrogen atom (N²⁴) which are the effective positions for chelation because they have high electronegative charge. From Figure 6 shows that the HOMO level is naturally positioned on the toluene moiety and azo-linkage’s nitrogen atoms (N^11&12) in addition to atoms of phenyl ring indicating that the O²² is the preferred position for nucleophilic attacking with the cations inductive that this atom has large coefficients of HOMO density and they are directed to the metallic centers [109]. The compounds separation energy values (ΔE = E_LUMO−E_HOMO) verified that the complexes (2–12) have ΔE ranged from 0.0593 to 0.1003 which are less stable and more reactive than the azo-Schiff (H₂L, 1) with ΔE = 0.1017 [110]. The term “electron affinity” refers to the energy expended if the system accepts an additional electron (EA = -E_LUMO) while the term “ionization potential” refers to the amount of energy needed to expel one electron from a species (IP = -E_HOMO) [109,111]. The absolute softness (σ = 1/η) and absolute hardness (η = (IP-EA)/2) could be utilized as a scale to reflect an atom's capacity to accept electrons and resist charge transfer, respectively [112]. The stability and reactivity of molecules might be predicted using these two factors [112]. Hard molecules have a large energy gap compared to soft molecules’ modest. As a result, the soft molecules’ reactivity is higher than that of hard molecules because soft molecules may easily transfer electrons to a receiver. The ligand molecules are the most efficient for coordination because, in coordinative compounds, the metal cations act as Lewis’ acids whereas the ligand molecules function as Lewis’ bases. The HOMO–LUMO diagram, which shows how reactive a molecule is, also shows a molecule's ability to transmit a charge and an electron to an accepting species. The molecule’s bioactivity is due to this property, which makes it easy for electrons to transfer from the HOMO to LUMO. The HOMO electron density is more evenly distributed in the backbone of complexes, as seen in Figure 7, which also denotes increased reactivity [113]. Consequently, the metallic cations are effectively chelated by the chelator with an appropriate σ value. This conclusion is assured from the chelator’s chemical potential value (Pi = -χ) [108,109]. From values of absolute electronegativities (χ = (IP + EA)/2), index of reactivity (ΔN_max= −Pi/η), chemical potentials (Pi = -χ), index of electrophilicity (ω = Pi ²/(2η)) [114], we can conclude that (i) the N_max is the measuring of the system's energy stabilization if it gains an extra electronic charge from its surroundings (ii) The chemical potential (Pi) recognizes the quantity of charge and the direction in which it is transferred, where the electrophiles are the chemical that may take electrons from their surrounding and must reduce their energy when doing. So, its Pi values must be negative as reinforced by the data in Table 3. (iii) The ω determines the ability of chemical species to receive electrons. Therefore, the high ω value indicates a well- functioning electrophile while low value implies a well-functioning nucleophile. (iv) The metallic cations were linked by azo-Schiff base (H₂L, 1) by oxygen atom of hydroxy group (O²²) and nitrogen atom of azomethine group (N²⁴) which own high electronegative charge demonstrating that these nitrogen and oxygen atoms are the effective sites for attachment. (v) The complexes dipole moment values are ranged from 3.68 to 15.4 Debye which are greater than that of azo-Schiff base, (H₂L, 1) (3.29 Debye), confirming that the attachment occurs by oxygen atom of hydroxy group (O²²) and nitrogen atom of azomethine group (N²⁴). Other energetic properties of the synthesized compounds which computed by DFT method are itemized in Table 4. The data of Table 4 showed that the calculated complexes’ binding energy are more than that of azo-Schiff base. This suggests that the stability of complexes is more than that of azo-Schiff base [115].

Figure 6. The HUMO-LUMO level 3D orbital images for the azo-Schiff base (H₂L).

Figure 7. The azo-Schiff base's molecular electrostatic potential (MEP).

3.2.3. Milliken atomic charges for the ligand

Milliken population analysis was used to determine the Milliken atomic charges of H₂L atoms by applying the DFT/B3LYP levels with basis set (6-311 + G(d,p)) in gas phase, which is shown in Figure S21 and listed in Table S4. The atomic charge playing a key role in the estimation of molecular system’s quantum chemical parameters as the atomic charge effect on the vibrational spectra, dipole moment, polarizability, electronic structure, and other properties of molecular system [58]. The Milliken analysis data discovered that the hydrogen atoms have positive charge (1.567 × 10⁻¹–3.742 × 10⁻¹ a.u), because of their losing electronic charge to the neighboring carbon atoms, whereas the hydrogen atoms attached to electronegative atoms (H-32 & H-33) have the highest positive charge of all hydrogen atoms because of their attachment to high electronegatively atoms (N,O) [58]. All oxygen and nitrogen atoms possess a negative Milliken atomic charge demonstrating that the oxygen and nitrogen atoms acted as electron receptors and server as donating atom for chelation.

3.2.4. Molecular electrostatic potential (MEP)

The charge distributions of azo-Schiff base and its variable charged zones are described and shown by MEP, which is estimated based on DFT calculation (Figure 7). MEP is typically employed as a reactivity diagram and to discover hydrogen bonding interactions, as well as nucleophilic and electrophilic reactions, in organic compounds [58,116]. The energy interaction between the molecule's nucleus protons, and electrons is represented by the MEP, V(r) at r(x, y, z) coordinates. The value of V(r) for the system under study was calculated using the following equation [58,116]. $\mathit{V}\left(\mathit{r}\right)=\sum _{\mathit{A}}\frac{{\mathit{Z}}_{\mathit{A}}}{{\mathit{R}}_{\mathit{A}}-\mathit{r}}\int \frac{\left(\mathit{r}\right)}{\left(\mathit{r}-\mathit{r}\right)}\mathit{d}\left(\mathit{r}\right)$where, r’, Z_A, and ρ(r′) indicates to the dummy integrated, and variables nucleus’ (A) charge that placed at R_A, electron charge density of molecule respectively [117]. The MEP employs a variety of colors to depict the various electrostatic potential values on the surface of molecule (from blue to red). Electron-rich locations are symbolized by a red color, whereas electro-positive locations are symbolized by the blue color, whereas electrostatic potentials close to zero are symbolized by the green color. The dihydrazone-oxime’s MEP 3D picture was displayed in Figure 7. As displayed in Figure 7, the greatest negative potential sites were located on the electronegatively atoms (O&N) which are the preferred locations for metallic chelation. The strongest positively potential regions (blue) were focused on a few carbon and hydrogen atoms, (+0.035 × 10⁻¹ to 3.742 × 10⁻¹ a.u). This result implies that nitrogen and oxygen atoms are more repellent than hydrogen atoms, which have more attractive characteristics.

3.3. In vitro anti-microbial screening

The antimicrobial assignment results of the ligand (H₂L,1), and its metal chelates (2–13) versus B. subtilis, E. coli (bacteria) and A. niger (fungus) is shown in Table 5 and shown in Figure 8. The complexes exhibited antifungal effects with inhibition zones ranging from 8 to 20 mm with indexes of activity from 47% to 118%. Both Zn²⁺ and UO₂²⁺ complexes exhibited the high potency antifungal activity with zones of inhibition equal to 20 and 17 mm with activity indexes 118% and 100% (of Am B), respectively. Ru³⁺ complex showed the least antifungal activity among all metal complexes investigated. The complexes exhibited bactericidal activity versus E. coli and B. subtilis with zones of inhibition in the 12–33- and 10–28-mm ranges with activity indexes in the 33%–92% and 26%–72%, respectively. The most potent complexes against E. coli are VO²⁺ and Ru³⁺ complex with inhibition zone equal to 33 mm with activity index 92% while versus B. subtilis the UO₂²⁺ is the most potent complex. The least potent complexes are Ni²⁺ and Mn²⁺ complexes. The complexes’ antimicrobial action is higher than that of the azo-Schiff base. This can be justified by Tweedy's chelation theory [118] and/or the concept of cell permeability [119], which can both be used to explain how complexes function. According to this theory, the polarity of metallic ions is diminished because of the sharing of their charge with the chelators’ donor atoms and the probability of π-electron delocalization over the chelate rings. This polarity diminishing leads to an increase in the lipophilic nature and this assists the penetration of complex molecules through the lipid membrane of the microbe cells. The overlap between the orbitals of the metal ions and the ligand’s molecular orbitals may be the cause of some complexes’ significant antibacterial activity. This overlapping may also help further distribute the positive charge of metal ions among the ligand molecule's donor groups. the resonance of π-electrons and the lipo-solubility of the whole chelate are increased and as a result, the metallic ions’ polarity is reduced. Therefore, it facilitates the diffusion of the complexes into microorganisms. The low activity of azo-Schiff base (H₂L, 1) and some of its metallic complexes against some microscopic strains might be attributed to the high polarity and hydrophobic nature of this compounds which decrease in the interaction with lipid membrane of the microorganism. Furthermore, these microbial strains’ efflux pumps prevent these compounds from penetrating the lipid membrane, providing protection against these complexes [120]. While chelation does indeed augment the biological activity of these complexes, several other factors come into play. These factors include the dipole moment, the geometry of the complexes, their concentration, molecular size, bond length, coordination sites, and hydrophobicity, all of which collectively influence their antimicrobial activity. Therefore, it is evident that the enhancement of antimicrobial activity is not solely attributed to chelation but rather results from a combination of various contributing factors [121,122]. The most exciting point of microbicides assignments is that this azo-Schiff base is inactive while other types of Schiff base such as ethyl-6-ethyl-2-((2- hydroxybenzylidene)amino)-4,5,6,7-tetrahydrothieno[2,3-c]pyridine-3-carboxylate [2], N-(4-fluorobenzylidene)-2-(quinolin-2-yloxy)acetohydrazide [123], 3-(-(1,5-dimethyl-3-oxo-2-phenyl-2,3-dihydro-1H-pyrazol-4-yl)methylene)hydrazonoindolin-2-one [124,125], 3-((hydroxyimino)butan-2-ylidene)benzohydrazide [126], Also the azo-compounds were more active than Azo-Schiff base [8,33,118,127], Also, their metal complexes show enhanced biological activity over azoSchiff base complexes. The lower activity of azoSchiff base and its complexes against both bacteria and fungi might be attributed to the presence of the azo group.

Figure 8. The microbicidal activity’s order of azo-Schiff base and its complexes (1 mg/mL/DMSO, at 28/37°C, pH 7.4).

Table 5. The prepared compounds’ microbicidal activity (mg/mL/DMSO, at 28/37°C, pH 7.4).

	Zone of inhibition (IZ, mm) and Index of activity (AI, %)
	A. niger		E. coli		B. subtilis
Compounds	IZ (mm)	AI (%)	IZ(mm)	AI (%)	IZ(mm)	AI (%)
DMSO	0	0	0	0	0	0
AmB	17	100	0	0	0	0
Tc	0	0	36	100	39	100
Ligand	0	0	0	0	0	0
VO^2+-Complex (2)	14	82	32	92	17	44
Fe^3+-Complex (3)	12	71	11	44	11	36
Ru^3+Complex (4)	8	47	33	92	16	41
Co^2+-Complex (5)	15	88	25	56	13	33
Mn^2+-Complex (6)	16	94	22	61	10	26
Ni^2+-Complex (7)	14	82	12	33	22	56
Cu^2+-Complex (8)	10	59	23	64	23	59
Cu^2+-Complex (9)	11	65	28	78	27	69
Cu^3+Complex (10)	13	76	20	56	24	62
Cu^2+-Complex (11)	12	71	20	56	24	62
Zn^2+-Complex (12)	20	118	32	89	25	64
UO2^2+-Complex (13)	17	100	28	78	28	72

4. Conclusion

Condensation of 2-hydroxy-5-(o-tolyldiazenyl)benzaldehyde with 1H-1,2,4-triazol-5-amine yielded a potentially bi-dentated Azo-Schiff base namely, 2-(((1H-1,2,4-triazol-5-yl)imino)methyl)-5-((4-nitrophenyl)diazenyl)phenol (H₂L, 1) which formed Ru³⁺, Fe³⁺, VO²⁺, Ni²⁺, Co²⁺, Mn²⁺, Cu², Zn²⁺ and UO₂²⁺ complexes of the types [M(HL)XYZ(H₂O)].nH₂O, [Cu(H₂L)SO₄(H₂O)] and [VO(H₂L)SO₄(H₂O)₂] in presence of counter ion (acetate, chloride, nitrate or sulphate). The structure of these prepared complexes was elucidated by physico-chemical, magnetic susceptibility, molar conductance, thermal studies, electronic, mass, nuclear magnetic resonance, and infrared spectra as well as theoretical studies. The results demonstrated that the metal-chelator stoichiometry in all complexes is 1M:1L in which the azo-Schiff base chelator worked as a monobasic/neutral NO-bidentate chelator bonding the metallic center via oxygen atom of deprotonated/protonated hydroxyl group and azomethine nitrogen atom adopting distorted octahedral, tetrahedral, or square planar geometry around the metallic center. Thermal gravimetric analysis (TGA) was used to examine thermal behavior, which focuses on the complexes’ ranges of stability and breakdown. Calculations based on DFT were employed to hypothetically analyze both the free azo-Schiff base’s 3D optimized structure and the metal complexes’ structures. The HOMO and LUMO energies were utilized for estimation the quantum chemical characteristics such as the energy gap, absolute electronegativity, chemical potential chemical hardness and chemical softness. The data of theoretical calculation displayed that all synthetics have negative HOMO and LUMO energies inductive to their stabilities. The in vitro antimicrobial tests reveal that the complexes have a high level of effectiveness as antifungal and antibacterial reagents more than the un-complexed azo-Schiff base ligand. Among the newly prepared complexes the Zn²⁺ and UO₂²⁺ complexes were the most effective complexes against fungi. They showed activity more than or resemble AmB. While the VO²⁺, Ru³⁺ and Zn²⁺ complexes were the most effective complexes versus bacteria with activity close to Tc. This finding suggests that a unique model of molecules with such high potency has a lot of promising. More investigation is necessary to find novel, potent congeners to address the microbial resistance problems with antimicrobial medicines. In our forthcoming research endeavors, our primary emphasis will be on the synthesis of innovative hybrid metal complexes incorporating azo-azomethine which is able to enhance the stability, decrease the toxicity, and enhance the biological activity, Furthermore, searching for novel synthetics is a valued target to overawe the microbial resistance problems of antimicrobial agents.

Disclosure statement

No potential conflict of interest was reported by the author(s).

Additional information

Funding

This work was supported by University of Jeddah [grant number UJ-22-DR-72].