# Nano-Azo Ligand and Its Superhydrophobic Complexes: Synthesis, Characterization, DFT, Contact Angle, Molecular Docking, and Antimicrobial Studies

Authors: Gehad G. Mohamed; Walaa H. Mahmoud; Ahmed M. Refaat
Journal: Journal of Chemistry
 (2020)

---

## Abstract

Metal complexes of the 2,2'-(1,3-phenylenebis(diazene-2,1-diyl))bis(4-aminobenzoic acid) diazo ligand (H2L) derived from m-phenylenediamine and p-aminobenzoic acid were synthesized and characterized by different spectral, thermal, and analytical tools. The H2L ligand reacted with the metal ions Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) as 1 : 1 stoichiometry. All complexes displayed an octahedral geometry according to the electronic and magnetic moment measurements. The IR spectra revealed the binding of the azo ligand to the metal ions via two azo nitrogen atoms and protonated carboxylate O in a neutral tetradentate manner. Both IR and 1H NMR spectra documented the involvement of the carboxylate group without proton displacement. The thermal studies pointed out that the complexes had higher thermal stability comparable with that of the free ligand. SEM images revealed the presence of the diazo ligand and its Cd(II) complex in a nanostructure form. The contact angle measurements proved that the Cd(II) complex can be considered as a superhydrophobic material. The molecular and electronic structure of H2L and [Cd(H2L)Cl2].H2O were optimized theoretically, and the quantum chemical parameters were calculated. The biological activities of the ligand, as well as its metal complexes, have been tested in vitro against some bacteria and fungi species. The results showed that all the tested compounds have significant biological activities with different sensitivity levels. The binding between H2L and its Cd(II) complex with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88) was predicted and given in detail using molecular docking.


---

## Body

## 1. Introduction

Azo compounds were highly colored compounds. They had widespread use as dyes and pigments in a variety of applications including fibers, coloring a large variety of leather, clothing, food, toys, medical devices, plastics, cosmetics, and dyeing of textile, as well as nonlinear and photo electronics, especially in optical information storage, biological medical studies, and advanced applications in organic synthesis [1–3]. In the development of metal complexes, the design and synthesis of a ligand was the most important step. They exhibited unique properties and novel reactivity due to electron-donor, electron-acceptor properties, structural, functional groups, and the position of the ligand in the coordination sphere [4, 5]. Azo dyes with hydrophobic characters have wide applications especially when these dyes exhibit superhydrophobic characters. Some applications of water repellency include clothing that will be both breathable and water repellant, umbrellas that stayed completely dry, building materials, paints, epoxies, and silicones. Superhydrophobic paints and epoxies could greatly reduce the cost of transporting goods, also eliminate many of the effects of ice storms and aircraft icing. Low-permeability paint, used as a water and vapor barrier, is the standard way of protecting metal surfaces from corrosion. Superhydrophobic biomaterials are also being evaluated in more demanding biological applications, such as the prevention of blood coagulation and drug delivery [6, 7].New diazo ligand (H2L) has the IUPAC name 2,2'-(1,3-phenylenebis(diazene-2,1-diyl))bis(4-aminobenzoic acid) (Scheme 1). This piece of work had devoted with the aim to synthesize Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) complexes with the new diazo ligand and to examine their physical properties involving spectral behaviors and the electrical conductance values, and to determine the efficiency of the synthesized complexes against pathogenic bacteria. The molecular structure and molecular docking were also carried out in order to illustrate the way of bonding of the azo ligand and its complexes with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88). In view of these findings, the contact angle measurements were performed to study the hydrophobic characters of the prepared compounds.Scheme 1
Pathway of formation of the H2L diazo ligand.


## 2. Experimental

### 2.1. Materials and Measurements

#### 2.1.1. Materials and Reagents

The chemicals used were of highest purity available and pure grade. Most of them were used without further purification.m-Phenylenediamine (Sigma), p-aminobenzoic acid (VEB Berlin-Chemie), CrCl3.6H2O (Sigma), MnCl2.2H2O (Sigma), NiCl2.6H2O (BDH), FeCl3.6H2O (Sigma), CoCl2.6H2O (Aldrich), CuCl2.2H2O (Merck), ZnCl2 (Strem Chemicals), and CdCl2 (Aldrich) were used. Organic solvents were spectroscopic pure from BDH and included ethanol and N,N-dimethylformamide. Hydrochloric acid, sodium nitrite, and sodium acetate (AR) were used.


#### 2.1.2. Solutions

1 × 10−3 M stock solutions of complexes were prepared by dissolving the appropriate amount of the complexes in DMF. It was used for conductivity measurement. For UV-Vis spectra measurement, 1 × 10−4 or 1 × 10−5 M solutions of the diazo ligand and metal complexes were prepared by accurate dilution from the previously prepared stock solutions.


#### 2.1.3. Measurements

Electronic and1H NMR spectra were carried out at room temperature using Shimadzu 3101PC spectrophotometer and 300 MHz Varian-Oxford Mercury, respectively. For 1H NMR spectra, solution of DMSO-d6 was used, and TMS was used as an internal standard. FT-IR spectra were recorded on a PerkinElmer 1650 spectrometer (4000–400 cm−1) in KBr pellets. Electron spin resonance spectra were also recorded on the JES-FE2XG ESR spectrophotometer at Microanalytical Center, Tanta University. Microanalyses of carbon, hydrogen, and nitrogen were carried out at Microanalytical Center, Cairo University, Egypt, using CHNS-932 (LECO) Vario Elemental Analyzer. Analyses of the metals were conducted by dissolving the solid complexes in concentrated HNO3 and dissolving the residue in deionized water. The metal content was carried out using inductively coupled plasma atomic absorption spectrometry (ICP-AES), Egyptian Petroleum Research Institute. Mass spectra were recorded by the EI technique at 70 eV using the MS-5988 GS-MS Hewlett–Packard instrument at Microanalytical Center, National Center for Research, Egypt. The molar magnetic susceptibility was measured on powdered samples using the Faraday method. The diamagnetic corrections were made by Pascal’s constant, and Hg[Co(SCN)4] was used as a calibrant. Molar conductivities of 10−3 M solutions of the solid complexes in DMF were measured using the Jenway 4010 conductivity meter. The thermogravimetric analyses (TG and DTG) of the solid complexes were carried out from room temperature to 800°C using a Shimadzu TG-50H thermal analyzer. The X-ray powder diffraction analyses were carried out by using Philips Analytical X-ray BV, diffractometer-type PW 1840. Radiation was provided by the copper target (Cu anode 2000 W) high-intensity X-ray tube operated at 40 KV and 25 mA, and SEM (scanning electron microscopy) for these samples was performed at the Egyptian Mineral Resources Authority, using SEM Model Quanta 250 FEG (field emission gun) attached with an EDX unit (energy-dispersive X-ray analyses), with accelerating voltage 30 KV, magnification 14x up to 1,000,000 and resolution for Gun.1n. Contact angle was measured by the contact-angle Attension Theta model (version 2.7), Biolin Scientific company, Finland, Egyptian Nanotechnology Center (EGNC). The heavy phase was water, where the volume of drop was 4 mL, and the light phase was air. Divergence and the receiving slits were 1 and 0.2, respectively. The antimicrobial activities were carried out at Microanalytical Center, Cairo University, Egypt.


### 2.2. Synthesis of 2,2’-(1,3-Phenylenebis(diazene-2,1-diyl))bis(4-aminobenzoic Acid) (H2L)

m-Phenylenediamine (2 g/0.01 mol) was dissolved in 50 ml of ethanol, while 10 ml concentrated hydrochloric acid was diluted with about 60 g of crushed ice; then, dropwise addition of m-phenylenediamine solution to the crushed ice was carried out. To this cold solution, sodium nitrite (5 g/20 mL water) was added and stirred for about 1 h till dark reddish colored solution was obtained. The coupling agent p-aminobenzoic acid (5.6 g/0.02 mol) was dissolved in 50 ml of ethanol and then was added to the cold mixture. The resulting solution was stirred well, and sodium acetate (3 g) was added to neutralization. The solid product was removed by filtration, washed with hot ethanol followed by diethyl ether, and dried in a vacuum desiccator over anhydrous calcium chloride. Preparation of the (H2L) ligand pathway is shown in Scheme 1.Yield 90%; m.p. > 300°C; reddish brown solid. Anal. Calcd for C20H16N6O4 (%): C, 59.20; H, 3.96; N, 20.70. Found (%): C, 58.74; H, 3.58; N, 20.17. IR (ν, cm−1): 3433br (OH), 1603m (N═N), 1660m (C═O carboxylic), 1543s (COOasym), 1384m (COOsym). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.25–7.91 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.12 ppm (s, 2H, COOH). UV-visible (λmax, nm): 375 (π–π∗), 271 (n–π∗).


### 2.3. Synthesis of Metal Complexes

H2L ligand (0.45 g/1.11 mmol) was dissolved in the mixture of 30 ml DMF. Then, the metal chloride salts (1.11 mmol) were dissolved in 10 ml ethanol, and their hot solutions were added to the clear hot solution of the diazo ligand. The precipitated complexes were filtered and collected. They were washed with a little amount of ethanol and dried over vacuum. The route of synthesis is shown in Scheme 2.Scheme 2
Synthesis of metal complexes of the diazo ligand (H2L).


#### 2.3.1. [CrH2LCl2]Cl.2H2O

Yield 92%; dark red; m.p. > 300°C. Anal. Calcd for Cr (C20H20Cl3CrN6O6) (%): C, 40.61; H, 3.34; N, 14.03; Cl, 17.79; Cr, 8.23. Found (%): C, 40.85; H, 3.05; N, 14.23; Cl, 17.46; Cr, 8.68. IR (ν, cm−1): 3423br (OH), 1610sh (N═N), 1655sh (C═O carboxylic), 1534m (COOasym), 1402w (COO sym), 537w (M─O), 426w (M─N). UV-visible (λmax, nm): 286 (π–π∗), 388 (n–π∗). Diffuse reflectance: 28,489, 24,780, and 13,449 cm−1 for 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P).


#### 2.3.2. [MnH2LCl2]

Yield 85%; pale green solid; m.p. > 300°C. Anal. Calcd for Mn (C20H16Cl2MnN6O4) (%): C, 44.20; H, 3.00; N, 15.80; Cl, 13.30; Mn, 10.49. Found (%): C, 44.54; H, 3.20; N, 15.34; Cl, 12.86; Mn, 10.37. IR (ν, cm−1): 3423br (OH), 1599m (N═N), 1655s (C═O carboxylic), 1553m (COOasym), 1389s (COOsym), 556w (M─O), 439w (M─N). UV-visible (λmax, nm): 269 (π–π∗), 374 (n–π∗). Diffuse reflectance: 26,645, 19,648, and 15,660 cm−1 for 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g (D) ⟶ 6A1g transitions.


#### 2.3.3. [FeH2LCl2]Cl.H2O

Yield 93%; black solid; m.p. > 300°C. Anal. Calcd for Fe(C20H18Cl3FeN6O5) (%): C, 41.02; H, 3.07; N, 14.63; Cl, 18.41; Fe, 9.32. Found (%): C, 41.27; H, 3.29; N, 14.35; Cl, 18.27; Fe, 9.57. IR (ν, cm−1): 3442br (OH), 1600sh (N═N), 1563s (C═O carboxylic), 1555m (COOasym), 1401s (COOsym), 578w (M─O), 442w (M─N). UV-visible (λmax, nm): 271 (π–π∗), 363 (n–π∗). Diffuse reflectance: 21,490 and 20,890 cm−1 for 6A1g ⟶ T2g (G), 6A1g ⟶ 5T1g and charge transfer, respectively.


#### 2.3.4. [CoH2LCl2]

Yield 91%; brown solid; m.p. > 300°C. Anal. Calcd for Co(C20H16Cl2CoN6O4) (%): C, 42.90; H, 2.99; N, 15.60; Cl, 13.20; Co, 11.53. Found (%): C, 43.12; H, 2.86; N, 15.11; Cl, 12.86; Co, 11.04. IR (ν, cm−1): 3425br (OH), 1612sh (N═N), 1652s (C═O carboxylic), 1537m (COOasym), 1384m (COOsym), 553w (M─O), 431w (M─N). UV-visible (λmax, nm): 260 (π–π∗), 343 (n–π∗). Diffuse reflectance: 23,110, 15,240, and 12,845 cm−1 for 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions.


#### 2.3.5. [NiH2LCl2]

Yield 93%; dark brown solid; m.p. > 300°C. Anal. Calcd for Ni(C20H16Cl2N6NiO4) (%): C, 43.90; H, 3.00; N, 15.71; Cl, 13.20; Ni, 11.45. Found (%): C, 43.51; H, 2.87; N, 15.22; Cl, 13.05; Ni, 11.04. IR (ν, cm−1): 3420br (OH), 1611 m (N═N), 1654s (C═O carboxylic), 1537m (COOasym), 1391s (COOsym), 541w (M─O), 432w (M─N). UV-visible (λmax, nm): 287 (π–π∗), 365 (n–π∗). Diffuse reflectance: 26,680, 14,757, and 13,122 cm−1 for ligand-metal charge transfer (LMCT), 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions.


#### 2.3.6. [CuH2LCl2].3/2H2O

Yield 91%; reddish brown solid; m.p. > 300°C. Anal. Calcd for Cu(C20H19Cl2CuN6O5.5) (%): C, 42.56; H, 3.76; N, 14.36; Cl, 12.91; Cu, 11.56. Found (%): C, 42.47; H, 3.35; N, 14.85; Cl, 12.55; Cu, 11.31. IR (ν, cm−1): 3450br (OH), 1614 m (N═N), 1650s (C═O carboxylic), 1554w (COOasym), 1392m (COOsym), 561w (M─O), 445w (M─N). UV-visible (λmax, nm): 270 (π–π∗), 346 (n–π∗). Diffuse reflectance: 24 729 and 13 290 cm−1 for 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions for LMCT.


#### 2.3.7. [Zn(H2L)Cl2].H2O

Yield 95%; reddish yellow solid; m.p. > 300°C. Anal. Calcd for Zn(C20H18Cl2N6O5Zn) (%): C, 42.97; H, 3.22; N, 15.04; Cl, 12.71; Cu, 11.96. Found (%): C, 42.81; H, 3.07; N, 14.88; Cl, 12.54; Zn, 11.54. IR (ν, cm−1): 3433br (OH), 1628 m (N═N), 1656s (C═O carboxylic), 1551m (COOasym), 1389s (COOsym), 573w (M─O), 429w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.20–7.89 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.40 ppm (s, 2H, COOH); UV-visible (λmax, nm): 280 (π–π∗), 395 (n–π∗).


#### 2.3.8. [Cd(H2L)Cl2].H2O

Yield 94%; orange solid; m.p. > 300°C. Anal. Calcd for Cd(C20H18CdCl2N6O5) (%): C, 39.67; H, 2.97; N, 13.88; Cl, 11.73; Cu, 18.57. Found (%): C, 39.56; H, 2.71; N, 13.68; Cl, 11.64; Cd, 18.13. IR (ν, cm−1): 3450br (OH), 1618 m (N═N), 1656s (C═O carboxylic), 1538s (COOasym), 1390m (COOsym), 546w (M─O), 435w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.25–7.93 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.38 ppm (s, 2H, COOH); UV-visible (λmax, nm): 275(π–π∗), 350 (n–π∗).


### 2.4. Spectrophotometric Studies

The absorption spectra were recorded for 1 × 10−4 M solutions of the free diazo ligand and its metal complexes dissolved in DMF. The spectra were scanned within the wavelength range from 200 to 700 nm.


### 2.5. Pharmacology

#### 2.5.1. Antimicrobial Activity

The antimicrobial activity was carried out using the diffusion agar technique [7, 8]. LB agar media surfaces were inoculated with four investigated bacteria (Gram-positive bacteria: Bacillus subtilis and Staphylococcus aureus; Gram-negative bacteria: Neisseria gonorrhoeae and Escherichia coli) and one strain of fungi (Candida albicans) and then transferred to a saturated disk with a tested solution in the center of Petri dishes (agar plates). DMSO served as a control. All Petri dishes were incubated at 25°C for 48 h where clear or inhibition zones were detected around each disk. Antibacterial activity was calculated [9, 10] by subtracting the diameter of the inhibition zone resulting with dimethylformamide from that obtained in each case. Amikacin and ketokonazole were used as reference compounds for antibacterial and antifungal activities, respectively. All experiments were performed as triplicate, and data plotted were the mean value.


### 2.6. Computational Methodology

The electronic structure calculations of H2L and Zn(II) complex were carried out using the Gaussian09 suite of program [11]. They were fully optimized employing DFT-based B3LYP method along with the LANL2DZ basis set. In order to incorporate the effect of the solvent around the molecule, the TD-DFT method (along with the LANL2DZ basis set) was used to calculate the electronic absorption spectra of the ligand and its Zn(II) complex. The contribution of the molecular orbital to HOMO and LUMO was also calculated.


### 2.7. Molecular Docking

Crystal structure ofS. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88) were used in this study. Molecular docking studies were performed using MOE 2008 software [12] in order to find out the possible binding modes of the most active compounds against the above receptors. It is an interactive molecular graphics program for calculating and displaying feasible docking modes of a receptor and ligand and complex molecules. It necessitates the ligand and the receptor as the input in the PDB format. The amino acid chain was kept, and the water molecules and cocrystallized ligands were removed. The structure of complexes in the PDB file format was created by Gaussian03 software. Crystal structure of HSA (human serum albumin) (PDB ID: 5FUO), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88) were downloaded from the Protein Data Bank (http://www.rcsb.org./pdb).


## 2.1. Materials and Measurements

### 2.1.1. Materials and Reagents

The chemicals used were of highest purity available and pure grade. Most of them were used without further purification.m-Phenylenediamine (Sigma), p-aminobenzoic acid (VEB Berlin-Chemie), CrCl3.6H2O (Sigma), MnCl2.2H2O (Sigma), NiCl2.6H2O (BDH), FeCl3.6H2O (Sigma), CoCl2.6H2O (Aldrich), CuCl2.2H2O (Merck), ZnCl2 (Strem Chemicals), and CdCl2 (Aldrich) were used. Organic solvents were spectroscopic pure from BDH and included ethanol and N,N-dimethylformamide. Hydrochloric acid, sodium nitrite, and sodium acetate (AR) were used.


### 2.1.2. Solutions

1 × 10−3 M stock solutions of complexes were prepared by dissolving the appropriate amount of the complexes in DMF. It was used for conductivity measurement. For UV-Vis spectra measurement, 1 × 10−4 or 1 × 10−5 M solutions of the diazo ligand and metal complexes were prepared by accurate dilution from the previously prepared stock solutions.


### 2.1.3. Measurements

Electronic and1H NMR spectra were carried out at room temperature using Shimadzu 3101PC spectrophotometer and 300 MHz Varian-Oxford Mercury, respectively. For 1H NMR spectra, solution of DMSO-d6 was used, and TMS was used as an internal standard. FT-IR spectra were recorded on a PerkinElmer 1650 spectrometer (4000–400 cm−1) in KBr pellets. Electron spin resonance spectra were also recorded on the JES-FE2XG ESR spectrophotometer at Microanalytical Center, Tanta University. Microanalyses of carbon, hydrogen, and nitrogen were carried out at Microanalytical Center, Cairo University, Egypt, using CHNS-932 (LECO) Vario Elemental Analyzer. Analyses of the metals were conducted by dissolving the solid complexes in concentrated HNO3 and dissolving the residue in deionized water. The metal content was carried out using inductively coupled plasma atomic absorption spectrometry (ICP-AES), Egyptian Petroleum Research Institute. Mass spectra were recorded by the EI technique at 70 eV using the MS-5988 GS-MS Hewlett–Packard instrument at Microanalytical Center, National Center for Research, Egypt. The molar magnetic susceptibility was measured on powdered samples using the Faraday method. The diamagnetic corrections were made by Pascal’s constant, and Hg[Co(SCN)4] was used as a calibrant. Molar conductivities of 10−3 M solutions of the solid complexes in DMF were measured using the Jenway 4010 conductivity meter. The thermogravimetric analyses (TG and DTG) of the solid complexes were carried out from room temperature to 800°C using a Shimadzu TG-50H thermal analyzer. The X-ray powder diffraction analyses were carried out by using Philips Analytical X-ray BV, diffractometer-type PW 1840. Radiation was provided by the copper target (Cu anode 2000 W) high-intensity X-ray tube operated at 40 KV and 25 mA, and SEM (scanning electron microscopy) for these samples was performed at the Egyptian Mineral Resources Authority, using SEM Model Quanta 250 FEG (field emission gun) attached with an EDX unit (energy-dispersive X-ray analyses), with accelerating voltage 30 KV, magnification 14x up to 1,000,000 and resolution for Gun.1n. Contact angle was measured by the contact-angle Attension Theta model (version 2.7), Biolin Scientific company, Finland, Egyptian Nanotechnology Center (EGNC). The heavy phase was water, where the volume of drop was 4 mL, and the light phase was air. Divergence and the receiving slits were 1 and 0.2, respectively. The antimicrobial activities were carried out at Microanalytical Center, Cairo University, Egypt.


## 2.1.1. Materials and Reagents

The chemicals used were of highest purity available and pure grade. Most of them were used without further purification.m-Phenylenediamine (Sigma), p-aminobenzoic acid (VEB Berlin-Chemie), CrCl3.6H2O (Sigma), MnCl2.2H2O (Sigma), NiCl2.6H2O (BDH), FeCl3.6H2O (Sigma), CoCl2.6H2O (Aldrich), CuCl2.2H2O (Merck), ZnCl2 (Strem Chemicals), and CdCl2 (Aldrich) were used. Organic solvents were spectroscopic pure from BDH and included ethanol and N,N-dimethylformamide. Hydrochloric acid, sodium nitrite, and sodium acetate (AR) were used.


## 2.1.2. Solutions

1 × 10−3 M stock solutions of complexes were prepared by dissolving the appropriate amount of the complexes in DMF. It was used for conductivity measurement. For UV-Vis spectra measurement, 1 × 10−4 or 1 × 10−5 M solutions of the diazo ligand and metal complexes were prepared by accurate dilution from the previously prepared stock solutions.


## 2.1.3. Measurements

Electronic and1H NMR spectra were carried out at room temperature using Shimadzu 3101PC spectrophotometer and 300 MHz Varian-Oxford Mercury, respectively. For 1H NMR spectra, solution of DMSO-d6 was used, and TMS was used as an internal standard. FT-IR spectra were recorded on a PerkinElmer 1650 spectrometer (4000–400 cm−1) in KBr pellets. Electron spin resonance spectra were also recorded on the JES-FE2XG ESR spectrophotometer at Microanalytical Center, Tanta University. Microanalyses of carbon, hydrogen, and nitrogen were carried out at Microanalytical Center, Cairo University, Egypt, using CHNS-932 (LECO) Vario Elemental Analyzer. Analyses of the metals were conducted by dissolving the solid complexes in concentrated HNO3 and dissolving the residue in deionized water. The metal content was carried out using inductively coupled plasma atomic absorption spectrometry (ICP-AES), Egyptian Petroleum Research Institute. Mass spectra were recorded by the EI technique at 70 eV using the MS-5988 GS-MS Hewlett–Packard instrument at Microanalytical Center, National Center for Research, Egypt. The molar magnetic susceptibility was measured on powdered samples using the Faraday method. The diamagnetic corrections were made by Pascal’s constant, and Hg[Co(SCN)4] was used as a calibrant. Molar conductivities of 10−3 M solutions of the solid complexes in DMF were measured using the Jenway 4010 conductivity meter. The thermogravimetric analyses (TG and DTG) of the solid complexes were carried out from room temperature to 800°C using a Shimadzu TG-50H thermal analyzer. The X-ray powder diffraction analyses were carried out by using Philips Analytical X-ray BV, diffractometer-type PW 1840. Radiation was provided by the copper target (Cu anode 2000 W) high-intensity X-ray tube operated at 40 KV and 25 mA, and SEM (scanning electron microscopy) for these samples was performed at the Egyptian Mineral Resources Authority, using SEM Model Quanta 250 FEG (field emission gun) attached with an EDX unit (energy-dispersive X-ray analyses), with accelerating voltage 30 KV, magnification 14x up to 1,000,000 and resolution for Gun.1n. Contact angle was measured by the contact-angle Attension Theta model (version 2.7), Biolin Scientific company, Finland, Egyptian Nanotechnology Center (EGNC). The heavy phase was water, where the volume of drop was 4 mL, and the light phase was air. Divergence and the receiving slits were 1 and 0.2, respectively. The antimicrobial activities were carried out at Microanalytical Center, Cairo University, Egypt.


## 2.2. Synthesis of 2,2’-(1,3-Phenylenebis(diazene-2,1-diyl))bis(4-aminobenzoic Acid) (H2L)

m-Phenylenediamine (2 g/0.01 mol) was dissolved in 50 ml of ethanol, while 10 ml concentrated hydrochloric acid was diluted with about 60 g of crushed ice; then, dropwise addition of m-phenylenediamine solution to the crushed ice was carried out. To this cold solution, sodium nitrite (5 g/20 mL water) was added and stirred for about 1 h till dark reddish colored solution was obtained. The coupling agent p-aminobenzoic acid (5.6 g/0.02 mol) was dissolved in 50 ml of ethanol and then was added to the cold mixture. The resulting solution was stirred well, and sodium acetate (3 g) was added to neutralization. The solid product was removed by filtration, washed with hot ethanol followed by diethyl ether, and dried in a vacuum desiccator over anhydrous calcium chloride. Preparation of the (H2L) ligand pathway is shown in Scheme 1.Yield 90%; m.p. > 300°C; reddish brown solid. Anal. Calcd for C20H16N6O4 (%): C, 59.20; H, 3.96; N, 20.70. Found (%): C, 58.74; H, 3.58; N, 20.17. IR (ν, cm−1): 3433br (OH), 1603m (N═N), 1660m (C═O carboxylic), 1543s (COOasym), 1384m (COOsym). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.25–7.91 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.12 ppm (s, 2H, COOH). UV-visible (λmax, nm): 375 (π–π∗), 271 (n–π∗).


## 2.3. Synthesis of Metal Complexes

H2L ligand (0.45 g/1.11 mmol) was dissolved in the mixture of 30 ml DMF. Then, the metal chloride salts (1.11 mmol) were dissolved in 10 ml ethanol, and their hot solutions were added to the clear hot solution of the diazo ligand. The precipitated complexes were filtered and collected. They were washed with a little amount of ethanol and dried over vacuum. The route of synthesis is shown in Scheme 2.Scheme 2
Synthesis of metal complexes of the diazo ligand (H2L).


### 2.3.1. [CrH2LCl2]Cl.2H2O

Yield 92%; dark red; m.p. > 300°C. Anal. Calcd for Cr (C20H20Cl3CrN6O6) (%): C, 40.61; H, 3.34; N, 14.03; Cl, 17.79; Cr, 8.23. Found (%): C, 40.85; H, 3.05; N, 14.23; Cl, 17.46; Cr, 8.68. IR (ν, cm−1): 3423br (OH), 1610sh (N═N), 1655sh (C═O carboxylic), 1534m (COOasym), 1402w (COO sym), 537w (M─O), 426w (M─N). UV-visible (λmax, nm): 286 (π–π∗), 388 (n–π∗). Diffuse reflectance: 28,489, 24,780, and 13,449 cm−1 for 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P).


### 2.3.2. [MnH2LCl2]

Yield 85%; pale green solid; m.p. > 300°C. Anal. Calcd for Mn (C20H16Cl2MnN6O4) (%): C, 44.20; H, 3.00; N, 15.80; Cl, 13.30; Mn, 10.49. Found (%): C, 44.54; H, 3.20; N, 15.34; Cl, 12.86; Mn, 10.37. IR (ν, cm−1): 3423br (OH), 1599m (N═N), 1655s (C═O carboxylic), 1553m (COOasym), 1389s (COOsym), 556w (M─O), 439w (M─N). UV-visible (λmax, nm): 269 (π–π∗), 374 (n–π∗). Diffuse reflectance: 26,645, 19,648, and 15,660 cm−1 for 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g (D) ⟶ 6A1g transitions.


### 2.3.3. [FeH2LCl2]Cl.H2O

Yield 93%; black solid; m.p. > 300°C. Anal. Calcd for Fe(C20H18Cl3FeN6O5) (%): C, 41.02; H, 3.07; N, 14.63; Cl, 18.41; Fe, 9.32. Found (%): C, 41.27; H, 3.29; N, 14.35; Cl, 18.27; Fe, 9.57. IR (ν, cm−1): 3442br (OH), 1600sh (N═N), 1563s (C═O carboxylic), 1555m (COOasym), 1401s (COOsym), 578w (M─O), 442w (M─N). UV-visible (λmax, nm): 271 (π–π∗), 363 (n–π∗). Diffuse reflectance: 21,490 and 20,890 cm−1 for 6A1g ⟶ T2g (G), 6A1g ⟶ 5T1g and charge transfer, respectively.


### 2.3.4. [CoH2LCl2]

Yield 91%; brown solid; m.p. > 300°C. Anal. Calcd for Co(C20H16Cl2CoN6O4) (%): C, 42.90; H, 2.99; N, 15.60; Cl, 13.20; Co, 11.53. Found (%): C, 43.12; H, 2.86; N, 15.11; Cl, 12.86; Co, 11.04. IR (ν, cm−1): 3425br (OH), 1612sh (N═N), 1652s (C═O carboxylic), 1537m (COOasym), 1384m (COOsym), 553w (M─O), 431w (M─N). UV-visible (λmax, nm): 260 (π–π∗), 343 (n–π∗). Diffuse reflectance: 23,110, 15,240, and 12,845 cm−1 for 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions.


### 2.3.5. [NiH2LCl2]

Yield 93%; dark brown solid; m.p. > 300°C. Anal. Calcd for Ni(C20H16Cl2N6NiO4) (%): C, 43.90; H, 3.00; N, 15.71; Cl, 13.20; Ni, 11.45. Found (%): C, 43.51; H, 2.87; N, 15.22; Cl, 13.05; Ni, 11.04. IR (ν, cm−1): 3420br (OH), 1611 m (N═N), 1654s (C═O carboxylic), 1537m (COOasym), 1391s (COOsym), 541w (M─O), 432w (M─N). UV-visible (λmax, nm): 287 (π–π∗), 365 (n–π∗). Diffuse reflectance: 26,680, 14,757, and 13,122 cm−1 for ligand-metal charge transfer (LMCT), 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions.


### 2.3.6. [CuH2LCl2].3/2H2O

Yield 91%; reddish brown solid; m.p. > 300°C. Anal. Calcd for Cu(C20H19Cl2CuN6O5.5) (%): C, 42.56; H, 3.76; N, 14.36; Cl, 12.91; Cu, 11.56. Found (%): C, 42.47; H, 3.35; N, 14.85; Cl, 12.55; Cu, 11.31. IR (ν, cm−1): 3450br (OH), 1614 m (N═N), 1650s (C═O carboxylic), 1554w (COOasym), 1392m (COOsym), 561w (M─O), 445w (M─N). UV-visible (λmax, nm): 270 (π–π∗), 346 (n–π∗). Diffuse reflectance: 24 729 and 13 290 cm−1 for 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions for LMCT.


### 2.3.7. [Zn(H2L)Cl2].H2O

Yield 95%; reddish yellow solid; m.p. > 300°C. Anal. Calcd for Zn(C20H18Cl2N6O5Zn) (%): C, 42.97; H, 3.22; N, 15.04; Cl, 12.71; Cu, 11.96. Found (%): C, 42.81; H, 3.07; N, 14.88; Cl, 12.54; Zn, 11.54. IR (ν, cm−1): 3433br (OH), 1628 m (N═N), 1656s (C═O carboxylic), 1551m (COOasym), 1389s (COOsym), 573w (M─O), 429w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.20–7.89 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.40 ppm (s, 2H, COOH); UV-visible (λmax, nm): 280 (π–π∗), 395 (n–π∗).


### 2.3.8. [Cd(H2L)Cl2].H2O

Yield 94%; orange solid; m.p. > 300°C. Anal. Calcd for Cd(C20H18CdCl2N6O5) (%): C, 39.67; H, 2.97; N, 13.88; Cl, 11.73; Cu, 18.57. Found (%): C, 39.56; H, 2.71; N, 13.68; Cl, 11.64; Cd, 18.13. IR (ν, cm−1): 3450br (OH), 1618 m (N═N), 1656s (C═O carboxylic), 1538s (COOasym), 1390m (COOsym), 546w (M─O), 435w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.25–7.93 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.38 ppm (s, 2H, COOH); UV-visible (λmax, nm): 275(π–π∗), 350 (n–π∗).


## 2.3.1. [CrH2LCl2]Cl.2H2O

Yield 92%; dark red; m.p. > 300°C. Anal. Calcd for Cr (C20H20Cl3CrN6O6) (%): C, 40.61; H, 3.34; N, 14.03; Cl, 17.79; Cr, 8.23. Found (%): C, 40.85; H, 3.05; N, 14.23; Cl, 17.46; Cr, 8.68. IR (ν, cm−1): 3423br (OH), 1610sh (N═N), 1655sh (C═O carboxylic), 1534m (COOasym), 1402w (COO sym), 537w (M─O), 426w (M─N). UV-visible (λmax, nm): 286 (π–π∗), 388 (n–π∗). Diffuse reflectance: 28,489, 24,780, and 13,449 cm−1 for 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P).


## 2.3.2. [MnH2LCl2]

Yield 85%; pale green solid; m.p. > 300°C. Anal. Calcd for Mn (C20H16Cl2MnN6O4) (%): C, 44.20; H, 3.00; N, 15.80; Cl, 13.30; Mn, 10.49. Found (%): C, 44.54; H, 3.20; N, 15.34; Cl, 12.86; Mn, 10.37. IR (ν, cm−1): 3423br (OH), 1599m (N═N), 1655s (C═O carboxylic), 1553m (COOasym), 1389s (COOsym), 556w (M─O), 439w (M─N). UV-visible (λmax, nm): 269 (π–π∗), 374 (n–π∗). Diffuse reflectance: 26,645, 19,648, and 15,660 cm−1 for 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g (D) ⟶ 6A1g transitions.


## 2.3.3. [FeH2LCl2]Cl.H2O

Yield 93%; black solid; m.p. > 300°C. Anal. Calcd for Fe(C20H18Cl3FeN6O5) (%): C, 41.02; H, 3.07; N, 14.63; Cl, 18.41; Fe, 9.32. Found (%): C, 41.27; H, 3.29; N, 14.35; Cl, 18.27; Fe, 9.57. IR (ν, cm−1): 3442br (OH), 1600sh (N═N), 1563s (C═O carboxylic), 1555m (COOasym), 1401s (COOsym), 578w (M─O), 442w (M─N). UV-visible (λmax, nm): 271 (π–π∗), 363 (n–π∗). Diffuse reflectance: 21,490 and 20,890 cm−1 for 6A1g ⟶ T2g (G), 6A1g ⟶ 5T1g and charge transfer, respectively.


## 2.3.4. [CoH2LCl2]

Yield 91%; brown solid; m.p. > 300°C. Anal. Calcd for Co(C20H16Cl2CoN6O4) (%): C, 42.90; H, 2.99; N, 15.60; Cl, 13.20; Co, 11.53. Found (%): C, 43.12; H, 2.86; N, 15.11; Cl, 12.86; Co, 11.04. IR (ν, cm−1): 3425br (OH), 1612sh (N═N), 1652s (C═O carboxylic), 1537m (COOasym), 1384m (COOsym), 553w (M─O), 431w (M─N). UV-visible (λmax, nm): 260 (π–π∗), 343 (n–π∗). Diffuse reflectance: 23,110, 15,240, and 12,845 cm−1 for 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions.


## 2.3.5. [NiH2LCl2]

Yield 93%; dark brown solid; m.p. > 300°C. Anal. Calcd for Ni(C20H16Cl2N6NiO4) (%): C, 43.90; H, 3.00; N, 15.71; Cl, 13.20; Ni, 11.45. Found (%): C, 43.51; H, 2.87; N, 15.22; Cl, 13.05; Ni, 11.04. IR (ν, cm−1): 3420br (OH), 1611 m (N═N), 1654s (C═O carboxylic), 1537m (COOasym), 1391s (COOsym), 541w (M─O), 432w (M─N). UV-visible (λmax, nm): 287 (π–π∗), 365 (n–π∗). Diffuse reflectance: 26,680, 14,757, and 13,122 cm−1 for ligand-metal charge transfer (LMCT), 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions.


## 2.3.6. [CuH2LCl2].3/2H2O

Yield 91%; reddish brown solid; m.p. > 300°C. Anal. Calcd for Cu(C20H19Cl2CuN6O5.5) (%): C, 42.56; H, 3.76; N, 14.36; Cl, 12.91; Cu, 11.56. Found (%): C, 42.47; H, 3.35; N, 14.85; Cl, 12.55; Cu, 11.31. IR (ν, cm−1): 3450br (OH), 1614 m (N═N), 1650s (C═O carboxylic), 1554w (COOasym), 1392m (COOsym), 561w (M─O), 445w (M─N). UV-visible (λmax, nm): 270 (π–π∗), 346 (n–π∗). Diffuse reflectance: 24 729 and 13 290 cm−1 for 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions for LMCT.


## 2.3.7. [Zn(H2L)Cl2].H2O

Yield 95%; reddish yellow solid; m.p. > 300°C. Anal. Calcd for Zn(C20H18Cl2N6O5Zn) (%): C, 42.97; H, 3.22; N, 15.04; Cl, 12.71; Cu, 11.96. Found (%): C, 42.81; H, 3.07; N, 14.88; Cl, 12.54; Zn, 11.54. IR (ν, cm−1): 3433br (OH), 1628 m (N═N), 1656s (C═O carboxylic), 1551m (COOasym), 1389s (COOsym), 573w (M─O), 429w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.20–7.89 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.40 ppm (s, 2H, COOH); UV-visible (λmax, nm): 280 (π–π∗), 395 (n–π∗).


## 2.3.8. [Cd(H2L)Cl2].H2O

Yield 94%; orange solid; m.p. > 300°C. Anal. Calcd for Cd(C20H18CdCl2N6O5) (%): C, 39.67; H, 2.97; N, 13.88; Cl, 11.73; Cu, 18.57. Found (%): C, 39.56; H, 2.71; N, 13.68; Cl, 11.64; Cd, 18.13. IR (ν, cm−1): 3450br (OH), 1618 m (N═N), 1656s (C═O carboxylic), 1538s (COOasym), 1390m (COOsym), 546w (M─O), 435w (M─N). 1H NMR (300 MHz, DMSO-d6, δ, ppm): 7.25–7.93 (m, 10H, Ar-H), 4.78 ppm (s, 4H, NH2) and 10.38 ppm (s, 2H, COOH); UV-visible (λmax, nm): 275(π–π∗), 350 (n–π∗).


## 2.4. Spectrophotometric Studies

The absorption spectra were recorded for 1 × 10−4 M solutions of the free diazo ligand and its metal complexes dissolved in DMF. The spectra were scanned within the wavelength range from 200 to 700 nm.


## 2.5. Pharmacology

### 2.5.1. Antimicrobial Activity

The antimicrobial activity was carried out using the diffusion agar technique [7, 8]. LB agar media surfaces were inoculated with four investigated bacteria (Gram-positive bacteria: Bacillus subtilis and Staphylococcus aureus; Gram-negative bacteria: Neisseria gonorrhoeae and Escherichia coli) and one strain of fungi (Candida albicans) and then transferred to a saturated disk with a tested solution in the center of Petri dishes (agar plates). DMSO served as a control. All Petri dishes were incubated at 25°C for 48 h where clear or inhibition zones were detected around each disk. Antibacterial activity was calculated [9, 10] by subtracting the diameter of the inhibition zone resulting with dimethylformamide from that obtained in each case. Amikacin and ketokonazole were used as reference compounds for antibacterial and antifungal activities, respectively. All experiments were performed as triplicate, and data plotted were the mean value.


## 2.5.1. Antimicrobial Activity

The antimicrobial activity was carried out using the diffusion agar technique [7, 8]. LB agar media surfaces were inoculated with four investigated bacteria (Gram-positive bacteria: Bacillus subtilis and Staphylococcus aureus; Gram-negative bacteria: Neisseria gonorrhoeae and Escherichia coli) and one strain of fungi (Candida albicans) and then transferred to a saturated disk with a tested solution in the center of Petri dishes (agar plates). DMSO served as a control. All Petri dishes were incubated at 25°C for 48 h where clear or inhibition zones were detected around each disk. Antibacterial activity was calculated [9, 10] by subtracting the diameter of the inhibition zone resulting with dimethylformamide from that obtained in each case. Amikacin and ketokonazole were used as reference compounds for antibacterial and antifungal activities, respectively. All experiments were performed as triplicate, and data plotted were the mean value.


## 2.6. Computational Methodology

The electronic structure calculations of H2L and Zn(II) complex were carried out using the Gaussian09 suite of program [11]. They were fully optimized employing DFT-based B3LYP method along with the LANL2DZ basis set. In order to incorporate the effect of the solvent around the molecule, the TD-DFT method (along with the LANL2DZ basis set) was used to calculate the electronic absorption spectra of the ligand and its Zn(II) complex. The contribution of the molecular orbital to HOMO and LUMO was also calculated.


## 2.7. Molecular Docking

Crystal structure ofS. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88) were used in this study. Molecular docking studies were performed using MOE 2008 software [12] in order to find out the possible binding modes of the most active compounds against the above receptors. It is an interactive molecular graphics program for calculating and displaying feasible docking modes of a receptor and ligand and complex molecules. It necessitates the ligand and the receptor as the input in the PDB format. The amino acid chain was kept, and the water molecules and cocrystallized ligands were removed. The structure of complexes in the PDB file format was created by Gaussian03 software. Crystal structure of HSA (human serum albumin) (PDB ID: 5FUO), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88) were downloaded from the Protein Data Bank (http://www.rcsb.org./pdb).


## 3. Results and Discussion

### 3.1. Characterization of the Free Diazo Ligand (H2L)

The synthesized diazo ligand (H2L) was soluble in DMF and DMSO solvents and of high melting point which indicated its stability. From the elemental analysis data of the ligand (Supplementary Table 1), it was found that the theoretical values were in agreement with the found values. The IR spectrum of the free diazo ligand was carried out in the range of 4000–400 cm−1, and the most effective bands are listed in Supplementary Table 2. The broad band observed at 3433 cm−1 can be attributed to the ν(O-H) stretching vibration. Also, band observed at 1660 cm−1 and sharp band at 1603 cm−1 can be attributed to C=O stretching vibration [13] and N=N [14], respectively. Also, the ν(COO)asym and ν(COO)sym stretching vibration bands were observed at 1543 and 1384 cm−1, respectively.1H NMR spectrum of the free diazo ligand (H2L) in DMSO-d6 exhibited the following signals: 7.25–7.91 ppm (m, 10H, Ar-H) [13], 4.78 ppm (s, 4H, NH2) [15–17], and 10.12 ppm (s, 2H, COOH). These signals indicated that three different types of protons were present in the diazo ligand. The position of NH2 was confirmed from carrying out the spectra in the deuterated solvent.The mass spectrum of the ligand (H2L) showed an accurate parent molecular ion peak at m/z 404 amu, matched with the theoretical molecular weight. Also, the mass spectrum displayed multiple peaks corresponding to successive degradation of the ligand. The base peak appeared at m/z = 134 due to the [C7H6NO2]+ ion. The other fragments gave the peaks at 104, 132, 136, 164, 240, and 268 amu with various intensities.


### 3.2. Characterization of Metal Complexes

#### 3.2.1. Elemental Analyses

The stoichiometry and formulation of the free diazo ligand (H2L) and its metal complexes were confirmed by their elemental analysis (Supplementary Table 1). The metal : ligand ratio was found to be 1 : 1 in all complexes, which has been arrived by estimating the carbon, hydrogen, nitrogen, chloride, and metal contents of the complexes. The elemental analyses of the ligand and its complexes revealed good agreement with the proposed structures. The ligand and its metal complexes have high melting points, and they were found to be air-stable. The diazo ligand was soluble in common organic solvents, and all the complexes were freely soluble in DMF and DMSO but insoluble in methanol, ethanol, and water. The results of elemental analyses suggested formulae [M(H2L)Cl2].nH2O (M = Co(II), Ni(II), Cu(II), Mn(II), Zn(II) and Cd(II)) and [M(H2L)Cl2]Cl.nH2O (M = Cr(III) and Fe(III)).


#### 3.2.2. Molar Conductance

The results given in Supplementary Table1 showed that the Cr(III) and Fe(III) complexes had molar conductivity values of 59 and 77 Ω−1 mol−1 cm2, respectively. It is obvious from these data that these chelates were ionic in nature, and they were electrolytes [18], but the molar conductivity values of Mn(II), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) chelates fall in the range of 11–31 Ω−1 mol−1 cm2 indicative of their nonelectrolytic nature [14].


#### 3.2.3. IR Spectral Studies

The infrared spectra of the H2L ligand and its metal complexes taken in the region 4000–400 cm−1 are listed in Supplementary Table 2 which gave the most useful assignments for bands diagnostic to the mode of coordination of the ligand. The band observed at 1603 cm−1 in the free diazo ligand (H2L) was assigned to the ν(N=N) stretching vibration. This band was found in the spectra of the complexes at 1610–1628 cm−1. Also, another band observed at 1660 cm−1 assigned to carbonyl of carboxylate ν(C=O) at the free diazo ligand (H2L) was found in the spectra of the complexes at 1649–1656 cm−1 (Supplementary Table 2). The shift in the azo groups of the diazo ligand confirmed the participation of the azo group in chelation.The broad band at 3433 cm−1 observed in the diazo ligand was shifted to 3423–3450 cm−1, which was attributed to the ν(O-H) stretching vibration. This shift can be attributed to the bond formation with metal ions and also the presence of hydrated water molecules. The broadening of this band may be attributed to the presence of the intramolecular hydrogen bond [19].Theνasym(COO−) and νsym(COO−) stretching vibrations were observed at 1543 and 1384 cm −1, respectively, for the free diazo ligand (H2L) [19–26]. Coordination with metal ions via the carboxylate O atom and nitrogen atom was indicated by the shift in the position of the νasym(COO) and νsym(COO) stretching vibration bands to 1534–1555 and1384–1402 cm−1, respectively, for ligand-metal complexes.New bands of low intensity were observed in the far-IR region in the range of 530–578 and 427–445 cm−1 which can be assigned to ν(M-O) and ν(M-N) stretching vibrations, respectively [20–28]. Therefore, from the IR spectral studies, it was concluded that the H2L ligand behaved as a neutral tetradentate ligand with N coordination sites and protonated carboxylate oxygen upon coordinated to the metal ions [20–27].


#### 3.2.4.1H NMR Spectral Studies

The1HNMR spectra of the diazo ligand and its Zn(II) complex were recorded in DMSO-d6 solution using tetramethylsilane (TMS) as the internal standard. 1H NMR spectrum of the ligand revealed the following signals at δ = 4.78 (m, 4H, NH2), 7.05–7.94 (m, 10H, Ar-H), and 10.18 ppm (s, 2H, COOH). The shift in the carboxylic acid protons (10.40 ppm) indicated that oxygen of (COOH) groups was coordinated to the Zn(II) ion. However, the singlet signal of the amino group that appeared at δ 4.78 ppm (s, 4H, NH2) means that the two singlet signals assigned to the two NH2 protons remained in their position as in the free diazo ligand (H2L), confirming the remaining of these groups inert towards coordination. The upfield shift in the spectrum of the complex relative to the free ligand can be attributed to the effect of metal ions in chelation.


#### 3.2.5. Mass Spectroscopy

The electronic mass spectrum of the [Zn(H2L)Cl2].H2O complex confirmed formation of this complex by showing a molecular ion peak at m/z = 540.81 g/mol corresponding to (M-H2O), which was very close to the expected molecular weight which was 540.50 g/mol, and it also showed a peak at m/z = 403 amu, respectively, corresponding to the ligand moiety. The intensities of these peaks showed the stability of the complex.


#### 3.2.6. Spectrophotometric Studies

UV-visible spectral data of the ligand and its chelate (10−4 M) solutions in DMF were recorded from 200 to 700 nm at room temperature using the same solvent as blank. The free ligand (H2L) exhibited broad bands in the UV-visible region at 375 and 271 nm. The first band was assigned to the intraligand n ⟶ π∗ electronic transition. The second broad band around 271 nm was assigned to the π ⟶ π∗ electronic transition. The n ⟶ π∗ and π ⟶ π∗ transition bands were shifted to 339–395 and 261–285 nm, respectively, indicating that the ligand coordinated to metal ions [29, 30].


#### 3.2.7. Magnetic Susceptibility and Electronic Spectral Studies

A comparison of the electronic spectra of the free diazo ligand (H2L) coordinated with those of the corresponding metal complexes was discussed. This can be considered as evidence for the complex formation. Additionally, the diffused reflectance spectra of metal complexes showed different bands at different wavelengths; each one was corresponding to a certain transition which suggested the geometry of the complexes. The diffused reflectance spectra of the complexes were dominated by intense intraligand charge-transfer bands.The diffused reﬂectance spectrum of the Cr(III) complex exhibited three bands at 28,489, 24,780, and 13,449 cm−1which may be assigned to the 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P) spin-allowed d–d transitions, respectively. The magnetic moment value was found to be 4.81 BM which indicated the presence of the Cr(III) complex in the octahedral geometry [31, 32].The diffused reﬂectance spectrum of the Mn(II) complex showed three bands at 26,645, 19,490, and 15.660 cm−1 assignable to 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g(D) ⟶ 6A1g transitions, respectively [31]. The magnetic moment value was found to be 5.54 BM which indicated the presence of the Mn(II) complex in the octahedral structure.From the diffused reﬂectance spectrum, it is observed that the Fe(III) chelate exhibited bands at 21,490 and 20,890 cm−1, which may be assigned to the 6A1g ⟶ T2g(G) and 6A1g ⟶ 5T1g transitions in the octahedral geometry of the complex [33]. The observed magnetic moment value of the Fe(III) complex was found to be 5.30 BM. The reflectance spectrum of the Co(II) complex displayed bands near 23,110, 15,240, and 12.845 cm−1, assigned to 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions, respectively. These bands suggested an octahedral geometry for the Co(II) complex. The magnetic moment value of the Co(II) complex was 5.16 BM which was consistent with the high-spin octahedral geometry.The diffused reflectance spectrum of the Ni(II) complex showed absorption bands at 20,680, 14,757, and 13,120 cm−1 which are attributed to 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions, respectively, and assignable to the octahedral geometry for the Ni(II) complex [34]. The magnetic moment value of the Ni(II) complex was 3.80 BM corresponding to the octahedral complex. The diffused reflectance spectrum of the Cu(II) complex exhibited broad bands in ranges 24,729 and 13,290 cm−1 assigned to 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions, respectively. These bands suggested distorted octahedral geometry for the Cu(II) complex. The magnetic moment value was found to be 1.78 BM indicating octahedral geometry [35].The Zn(II) and Cd(II) complexes were diamagnetic. According to the empirical formulae, an octahedral geometry was proposed for these chelates.


#### 3.2.8. ESR Studies of the Cu(II)-H2L Complex

The ESR spectrum of the Cu(II) complex was recorded in DMSO at 300 and 77 K (Figure1). The observed order for the Cu(II) complex (A║ (197) > A┴ (98.88); g║ (2.445) > g┴ (2.333)) indicated that the complex exerted an octahedral geometry [35–37]. The observed value of G for the Cu(II) complex (G = 1.421) implied that the exchange coupling was not present, and misalignment was appreciable. The trend g|| > g⊥ > ge (2.0023) showed that the unpaired electron was localized in the dx2-y2 orbital of the Cu(II) ion in the complex [35–38]. The giso (2.0023) value was less than 2.3 which indicated the covalent character of the metal-ligand bond, and the α2 value (0.76) suggested the presence of in-plane covalency. The calculated value (g║/A║) 124 cm for the complex was consistent with the slightly distorted structure. The orbital reduction factors K║ and K⊥ were estimated from the following expressions: K║ = (g║ −2.0023) ΔE/8λ and K⊥ = (g║ −2.0023) ΔE/2λ, λ = −828 cm−1 (spin-orbit coupling constant for the free ion), and K║ (0.718) > K⊥ (0.45) showed poor in-plane π bonding.Figure 1
ESR spectrum of the [Cu(H2L)Cl2].3/2H2O complex


#### 3.2.9. Powder X-Ray Diffraction and SEM

The XRD patterns indicated crystalline nature for the H2L diazo ligand and the Cd(II) complex only. It can be easily seen that the pattern of the H2L diazo ligand differed from its metal complexes, which may be attributed to the formation of a well-deﬁned distorted crystalline structure. Probably, this behavior was due to the incorporation of water molecules into the coordination sphere. On comparing the XRD spectra of the chelates with the XRD spectra of the free ligand, it was concluded that all chelates under study can be considered to have amorphous structures as they lack sharp peaks except the Cd(II) complex where it had a crystalline structure. Therefore, the nonsimilarity of the XRD pattern between the metal ions and chelates suggested that these chelates had different phase structures than the free diazo ligand (H2L). Such facts suggested that H2L and the Cd(II) complex were crystalline, while its Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), and Zn(II) complexes were amorphous.The average crystallite size (ξ) can be calculated from the XRD pattern according to the Debye–Scherrer equation [39, 40]:(1)ξ=Kλβ1/2cosθ.The equation uses the reference peak width at angleθ, where λ is the wavelength of the X-ray radiation (1.541874 Å), K is a constant taken as 0.95 for organic compounds [37], and β1/2 is the width at half maximum of the reference diffraction peak measured in radians. The dislocation density, δ, is the number of dislocation lines per unit area of the crystal. The value of δ is related to the average particle diameter (ξ) by the following relation [41, 42]:(2)δ=1ξ2.The value ofδ was calculated and found to be 33.78 and 26.77 nm, and the value of δ was 8.76 × 10−4 and 1.39 × 10−3 nm−2 for H2L and Cd(II) complex, respectively, giving evidence for the formation of the free azo ligand and its cadmium(II) complex in a nanosize.The SEM micrographs of H2L and its Cd(II) are given in Figure 2. There is a significant difference in the morphological structures, and the particle size is in the nanometer range. The particle size of the ligand was 64–67 nm, while the Cd(II) complex was 20–32 nm.Figure 2
The SEM micrographs of (a) Cd(II)-H2L and (b) H2L.


#### 3.2.10. Contact Angle

For two centuries or more of research, wetting phenomena still remain popular topics of investigations. Young’s static and dynamic contact angle equations [6, 7] are the most common and relevant parameters describing the wetting property of a surface with respect to liquid water:(3)cosθY=σsv−σs1σ1v,where θY = Young’s contact angle, σsv = surface tension (energy per unit surface) of the solid-vapor interface, σsl = surface tension of the solid-liquid interface, and σlv = surface tension of the liquid-vapor interface.The mean contact angles were measured for H2L and its metal complexes (Supplementary Figure 3). The data listed in Table 1 proved that the diazo ligand and all the complexes are hydrophobic, and the Cd(II) complex can be considered as a superhydrophobic material. Superhydrophobic surfaces are actively studied across a wide range of applications and industries and are now finding increased use in the biomedical areas. A partial list includes clothing that will be both breathable and water repellant, umbrellas that stayed completely dry, building materials, paints, epoxies, goods transport, and anticorrosion. Superhydrophobic materials such as this promising nano-Cd(II) complex exhibit a number of unique properties that arise from the high roughness of a low surface energy material that stabilizes a nonwetted state [6, 7].Table 1
The mean contact angle of H2L and its complexes.
CompoundMean contact angle (°)Wettability classificationSurface morphology[Cr(H2L)Cl2]Cl.2H2O137HydrophobicRough and porous[Mn(H2L)Cl2]128HydrophobicRough and porous[Fe(H2L)Cl2]Cl.H2O142HydrophobicRough and porous[Co(H2L)Cl2]123HydrophobicRough and porous[Ni(H2L)Cl2]132HydrophobicRough and porous[Cu(H2L)Cl2].3/2H2O133HydrophobicRough and porous[Zn(H2L)Cl2].H2O128HydrophobicRough and porous[Cd(H2L)Cl2].H2O153SuperhydrophobicVery rough and very porousH2L azo dye ligand103HydrophobicRough and porous


#### 3.2.11. Thermal Analyses

Thermogravimetric analyses of the metal complexes under investigation were used to get information about the thermal stability of these new complexes, to decide whether the water molecules (if present) were inside or outside the inner coordination sphere of the central metal ion, and to suggest a general scheme for thermal decomposition of these chelates. In the present investigation, heating rates were suitably controlled at 10°C min−1 under nitrogen atmosphere, and the weight loss was measured from the ambient temperature up to ∼1000°C. The TG curves of the complexes are listed in Table 2. The weight losses for each chelate were calculated within the corresponding temperature ranges.Table 2
Thermoanalytical results (TG and DTG) of the H2L ligand and its metal complexes.
ComplexTG range (oC)DTG maxn∗Mass loss estim (calcd)% (total mass loss)AssignmentResiduesDiazo ligand (H2L)30–355216136.53 (38.61)Loss of 4NH3 and 2CO2—355–1000476,602261.76 (61.38) 100.0 (99.90)Loss of C18H14N2[Cr(H2L)Cl2]Cl.2H2O45–1157918.88 (8.53)Loss of 2H2O and NH31/2 Cr2O3115–310210130.00 (29.90)Loss of 3/2Cl2, 2N2, and CH4310–990377,865241.30 (43.10) 80.18 (81.53)Loss of C17H12NO2.5[Mn(H2L)Cl2]30–435234158.98 (58.86)Loss of 2N2,Cl2, and C9H5N27C + MnO435–1000592120.88 (20.18) 79.66 (79.04)Loss of C4H11O3[Fe(H2L)Cl2]Cl.H2O45–1007513.29 (3.07)Loss of H2O5C + 1/2Fe2O3100–355244147.89 (48.33)Loss of 3/2Cl2, 3N2, 3CH4, and CO2355–1000485,565,803325.37 (24.63) 76.62 (76.03)Loss of C11H4O0.5[Co(H2L)Cl2]30–260177112.23 (12.45)Loss of 2NH3 and1/2Cl22C + CoO260–47536416.33 (6.64)Loss of 1/2 Cl2475–890640160.3 (60.03) 78.91 (79.09)Loss of C18H13N5O[Ni(H2L)Cl2]40–290199131.24 (31.27)Loss of 4NH3,N2, and Cl28C + NiO290–970395,883235.57 (36.70) 67.85 (67.95)Loss of C12H4O3[Cu(H2L)Cl2].3/2H2O45–1458814.71 (4.77)Loss of 3/2H2O2C + CuO145–990287177.67 (76.92) 82.38 (81.70)Loss of C18H16Cl2N6O3[Zn(H2L)Cl2].H2O45–1308615.42 (5.72)Loss of H2O and 1/2N23C + ZnO130–270205110.36 (10.26)Loss of 2N2270–1000348,467263.4 (63.20) 79.17 (79.18)Loss of C17H16Cl2NO3[Cd(H2L)Cl2].H2O40–1309812.95 (2.90)Loss of H2OCdO130–355222125.63 (25.62)Loss of 3N2 and Cl2355–1000414,591,833350.15 (50.25) 78.73 (78.77)Loss of C20H16O3n∗ = number of decomposition steps.The thermal decomposition process of H2L involved three decomposition steps. Decomposition of H2L started at 30°C and finished at 100°C with three stages. The first stage of decomposition involved the removal of four molecules of NH3 and two molecules of carbon dioxide gas in the 30–355°C temperature range and was accompanied by a weight loss of 36.53% (calcd. = 36.41%). The second and third stages of decomposition occurred in the 355–1000°C temperature range, corresponding to the loss of the C18H4N2 molecule and were accompanied by a weight loss of 62.76% (calcd. = 62.59%).Decomposition of the [Cr(H2L)Cl2]Cl.2H2O complex started at 45°C and ﬁnished at 1000°C with four stages at four maxima at 79, 210, 377, and 865°C and was accompanied by a weight loss of 80.29%, corresponding to the loss of 2H2O, 3/2Cl2, 2N2, CH4, and C17H12NO2.5 molecules which was very close to the calculated value of 81.53%.Thermogravimetric curve for the [Mn(H2L)Cl2] complex showed two weight loss events. The ﬁrst step of decomposition occurred within the range of 30–435°C, with a maximum temperature at 234°C, and it corresponds to the loss of Cl2, 2N2 gases and C9H5N2 molecules. The second step of decomposition occurred in the range of 435–1000°C with the maximum temperature at 595oC. There were signiﬁcant mass loss events due to the decomposition of the C4H11O3molecule (Table 2).The thermal decomposition of the [Fe(H2L)Cl2]Cl.H2O complex proceeded via five degradation steps. The ﬁrst step of decomposition occurred within the temperature range from 45 to 100°C, with a maximum temperature at 75°C, and it corresponds to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 100–355°C, with one maxima at 244°C, and corresponds to the elimination of 3/2Cl2, 3N2, 3CH4, and CO2 with a mass loss of 47.89% (calcd. = 48.33%). The last three steps of decomposition occurred within the temperature range of 355–1000°C with three maxima at 485, 565, and 803°C and were simultaneously decomposed to ferric oxide contaminated with carbon as a ﬁnal product. They were accompanied by a weight loss of 25.37% (calcd. = 24.63%) and corresponded to the loss of the C11H4O0.5 molecule. The actual weight loss from these five steps was 76.62% which was close to the calculated value of 76.03%.The thermal degradation of the [Co(H2L)Cl2] complex took place in three degradation stages within the temperature range from 30 to 990oC. The ﬁrst stage of decomposition occurred at maximum 177°C and was accompanied by a weight loss of 12.23% (calcd. = 12.45%) corresponding to the loss of 2NH3 and 1/2Cl2 gases. The second and third steps of decomposition occurred at two maxima 364 and 640°C and were accompanied by a weight loss of 6.33% (calcd. = 6.64%) and 60.30% (calcd. = 60.03%), respectively, corresponding to the loss of 1/2Cl2 gas and C18H3N5O molecule giving CoO contaminated with carbon as a ﬁnal product. Theoretically, the weight loss in these steps was 78.91% which agrees with the experimental value of 79.09%.The [Ni(H2L)Cl2] complex lost upon heating 4NH3, Cl2, and N2 gases in the ﬁrst step of decomposition within the temperature range of 40–290oC. The last two steps of decomposition occurred at two maxima at 395 and 880°C. These steps were associated with the loss of the C12H4O3 molecule forming NiO contaminated with carbon as a ﬁnal product. The actual total weight loss of 67.85% was in agreement with the calculated total weight loss value of 67.95%.The complex [Cu(H2L)Cl2].3/2H2O lost upon heating all water molecules at one maxima 90oC with an estimated weight loss of 4.70% (calcd. = 4.77%). The dehydrated Cu(II) complex was simultaneously decomposed to CuO and 2C at the 287oC maximum temperature. The total weight loss amounted to 82.38% (calcd. = 81.74%).The [Zn(H2L)Cl2].H2O complex was thermally decomposed into four steps within the temperature range from 45 to 1000oC. The ﬁrst decomposition step with an estimated mass loss of 5.42% (calcd. = 5.72%) occurred within the temperature range from 45 to 130oC. This step may be attributed to the liberation of the hydrated water molecules and NH3 gas. The remaining three decomposition steps of the Zn(II) complex were observed within the temperature range from 130 to 1000oC. Estimated mass losses of 73.76% (calcd. = 74.42%) were reasonably accounted for the removal of two 2N2 gases and a C17H16Cl2NO3 molecule in two steps.TG curve of the [Cd(H2L)Cl2].H2O complex displayed a five-step decomposition. The ﬁrst step of decomposition occurred within the temperature range from 40 to 130oC, with a maximum temperature at 98 oC, and corresponded to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 130–355 oC, with one maxima at 222oC, and corresponded to the elimination of Cl2 and 3N2 gases with a mass loss of 25.62% (calcd. = 25.63%). The last three steps of decomposition occurred within the temperature range of 355–1000oC with three maxima at 414, 591, and 833oC and were simultaneously decomposed to CdO as a ﬁnal product. They were accompanied by a weight loss of 50.25% (calcd. = 50.15%) corresponding to the loss of the C20H16O3 molecule. The weight loss from these five steps was 78.77% which was close to the calculated value of 78.73%.


### 3.3. Geometry Optimization

The geometric parameters (bond lengths and bond angles) of the optimized structures of the H2L azo ligand and its Cd(II) complex (Figure 3) are listed in Supplementary Table 5. From the analysis of these data, the following remarks were found:(1)
There was an elongation in the coordination bonds after complexation, and a large variation in N(11)-N(14), N(12)-N(13), O(35)-H(40), and O(38)-H(39) bond lengths of the ligand occurred upon complexation. It became slightly longer confirming that the coordination took place via azo N atoms and protonated O groups.(2)
All the active groups taking part in coordination have bonds longer than those already existed in the ligands (for example, N=N).(3)
The bond angles of the free ligand moiety were altered somewhat upon coordination; also, the bond angles surrounding the central cadmium atom in the complex lie in the range reported for octahedral geometry [43].(4)
The decrease in the metal-chloride angles may be attributed to the intramolecular hydrogen bond [43].Figure 3
The optimized structure of (a) H2L and (b) [Cd(H2L)Cl2].H2O.
(a)(b)


### 3.4. Molecular Electrostatic Potential (MEP)

The MEP is a plot of electrostatic potential mapped onto the constant electron density surface. It is also very useful in the research of the molecular structure with its physiochemical property relationship. It is also can be used to understand the molecular interactions and predict the reactive sites for electrophilic and nucleophilic attacks [43]. The MEP is directly related to chemical reactivity. Red color represented the regions of negative electrostatic potential. The negative regions are related to electrophilic reactivity. However, the electron-poor region has blue color (favor site for the nucleophilic attack) [43, 44], but the region with green color points to the neutral electrostatic potential region. Figure 4 represents 3D plots of the MEP for the parent ligand and its Cd(II) complex. It was found that oxygen and nitrogen atoms of free H2L were surrounded by a greater negative charge surface, making these sites potentially more favorable for the electrophilic attack (Figure 4(a)), where a greater negative charge was surrounded to the metal center (Figure 4(b)) confirming the complexation process.Figure 4
Molecular electrostatic potential map of (a) H2L and (b) [Cd(H2L)Cl2].H2O. The electron density isosurface is 0.004 a.u.
(a)(b)


### 3.5. Mulliken Charges

Supplementary Table6 and Supplementary Figure 6 clarify Mulliken atomic charges for some important atoms. The calculated charge value for N13 and N14 is −0.08 and −0.007 a.u., respectively, in contrast to the expected one (−1 a.u.), which indicated electron transfer from these atoms to the metal center. Furthermore, decrease of oxygen and chlorine formal charge from the expected values is a reason for migration of electrons to the cadmium center in order to forming the complex. Finally, this charge transfer from donor atoms to the metal center causes the calculated value to decrease to −0.756 a.u. against the formal charge of +2 for the cadmium atom [43, 44].


### 3.6. Molecular Parameters

Figure5 shows the molecular orbital representation of the H2L ligand and its Cd(II) complex along with their HOMO, LUMO energies and energy band gaps. Both the highest occupied molecular orbital (HOMO) and lowest unoccupied molecular orbital (LUMO) were the main orbitals that participate in chemical stability [11, 44, 45]. The HOMO represented the ability to donate an electron; LUMO as an electron acceptor represents the ability to obtain an electron. In other words, the energy of the HOMO is directly related to the ionization potential; LUMO energy is directly related to the electron affinity. Hence, these frontier energies play an important role in the electric and optical properties. The energy gap between the HOMO and LUMO energies is a critical parameter in determining molecular electrical transport properties due to providing measurement of electron conductivity. In addition, this energy gap characterizes the molecular stability and spectroscopic properties of the molecular systems. The smaller energy gap describes a chemically soft molecule which can be easily polarizable. The HOMO-LUMO energy gap has proved to be an important tool in determining the kinetic stability and chemical reactivity of a molecule. The lower energy gap is a suitable condition where a molecule can be exited easily, whereas a higher energy gap leads to higher kinetic stability but lower chemical reactivity of the molecule. The difference between the energy gap for the ligand reflects the presence of the complexation status. The increase in the global electrophilicity value attributed to the highest capacity of accepted electrons, so the ligand has a highly powerful donation ability. The calculation of both absolute hardness (η) and absolute softness (σ) parameters is useful to recognize the molecular stability and reactivity. The calculations of the binding energy indicated that the increase of the value of the calculated binding energy of the complex compared to that of the free ligand indicated that the stability of the formed Cd(II) complex was higher than that of the free diazo ligand.Figure 5
Surface phase of frontier orbitals of H2L and [Cd(H2L)Cl2].H2O.Additional parameters such as chemical potentialsPi, global electrophilicity ω, global softness S, electrophilicity index (χ), and additional electronic charge ∆Nmax were calculated for the free H2L ligand and the Cd-H2L complex (Table 3). The high ω value of both free H2L and Cd(II) complex suggested a great chance and priority for biological activity which is further confirmed by the experimental data.Table 3
The calculated quantum chemical parameters of H2L and its Cd(II) complex.
The quantum parameterH2L[Cd(H2L)Cl2].H2OE (a.u.)−1402.46−140.88Dipole moment (Debye)1.3718.022EHOMO (eV)−6.095−5.61ELUMO (eV)−2.977−4.096ΔE (eV)3.3491.508`χ(eV)4.5344.853η (eV)1.5590.757σ (eV)−10.6411.321Pi (eV)−4.534−4.853S (eV)−10.3210.661ω (eV)6.59315.559∆Nmax2.96.412The mentioned quantum chemical parameters were calculated with the help of the following equations [43, 44]:(4)ΔE=ELUMO−EHOMO,χ=−EHOMO+ELUMO2,η=ELUMO−EHOMO2,σ=1η,Pi=−χ,S=12η,ω=Pi22η,ΔNmax=Piη.


### 3.7. Vibrational Properties

In order to acquire the spectroscopic signature of the diazo ligand and its complexes, a frequency calculation analysis was performed. Vibrational properties had been investigated to determine vibrational modes connected with the molecular structure of the cadmium (II) complex using FT-IR spectra based on the optimized geometry. It is well known that the vibrational frequencies computed at quantum chemical methods such as DFT levels contain well-known systematic errors. The scaling factor of 0.96 for the LanL2DZ level was used to correct the effects of anharmonicity and neglected part of electron correlation [11, 44, 45]. The theoretical and the experimental spectra are shown in Supplementary Figure 8. It is also noteworthy that experimental results belong to the solid phase, while theoretical calculations belong to a gas phase.The stretching vibration of a free or nonhydrogen-bonded OH group of H2L appears at 3664 cm−1 [11, 45]. υ(COO−)asy and υ(COO−)sy stretching of the free ligand appeared at 1597 cm−1 and 1392 cm−1, respectively. The theoretical spectrum of the free azo ligand showed υNH2 and υN = N modes at 3637 and 1694 cm−1, respectively. In the Cd(II) complex, the chelating ring occurred by means of the ligand bonding to the cadmium (II) ion with a tetradentate N,O chelated mode. It was observed from Figures 6 that the wavenumber interval between the υ(COO−)asy and υ(COO −)sy stretching modes is an effective indicator of forming complexation. The wavenumber value of υN = N for the complex was lower than free H2L which was an indicator of the coordination through the nitrogen atom. The spectrum of the Cd(II) complex showed υNH2 at a wave number very close to free H2L at 3639 cm−1 confirming that this group did not participate in the complex formation process. υM-O, υM-N, and υM-Cl modes of the Cd(II) complex were found at 564, 487, and 326 cm−1, respectively. The calculated and experimental data were very close to each other (Supplementary Figure 8).Figure 6
The proposed structures of H2L and its complexes.


### 3.8. Structural Interpretation of Metal Complexes

The structures of the complexes of the H2L diazo ligand with Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) ions were conﬁrmed by the elemental analyses, IR, molar conductance, magnetic, mass, solid reﬂectance, and thermal analysis data. Therefore, from the IR spectra, it was concluded that H2L behaved as a neutral tetradentate ligand coordinated to the metal ions via two nitrogen atoms of azo groups (N=N) and two oxygen atoms of protonated carboxylate groups (COO). From the molar conductance data, it was found that the complexes were electrolytes. On the basis of the above observations and from the magnetic and solid reﬂectance measurements, octahedral geometry was suggested for the investigated complexes (Figure 6).


### 3.9. Antimicrobial Activity of Metal Complexes

#### 3.9.1. Antifungal Activities

The preliminary fungi toxicity screening of the H2L diazo ligand and complexes was performed against Candida albicans in vitro by the diffusion technique [46]. H2L free ligand and Cr(III), Mn(II), and Fe(III) complexes showed no fungal growth inhibition, but Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) complexes showed fungal growth inhibition at different rates (Figure 7). The Cd(II) and Ni(II) complexes in this study were nearly two times more active than the ketokonazole standard against the Candida albicans microorganism used (Table 4).Figure 7
Biological activity of H2L and its complexes.Table 4
Biological activity of the H2L ligand and its metal complexes.
SampleInhibition zone diameter (mm/mg sample)Bacillus subtilisEscherichia coliNeisseria gonorrhoeaeStaphylococcus aureusCandida albicansControl: DMSO0.00.00.00.00.0Ligand (H2L)0.00.00.00.00.0[Cr(H2L)Cl2]Cl.2H2O141515150.0[Mn(H2L)Cl2]91090.00.0[Fe(H2L)Cl2]Cl.H2O9990.00.0[Co(H2L) Cl2]1616151612[Ni(H2L)Cl2]1414151316[Cu(H2L)Cl2].3/2H2O1212121310[Zn(H2L)Cl2].H2O111213159[Cd(H2L)Cl2].H2O1815132317Amikacin69760Ketokonazole————9


#### 3.9.2. Antibacterial Activities

The antibacterial activities of the H2L diazo ligand and its complexes against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli are presented in Table 4. The H2L ligand was found to have no activity at all towards both different types of bacteria (Gram-positive Bacillus subtilis and Staphylococcus aureus and Gram-negative Neisseria gonorrhoeae and Escherichia coli). This can be attributed to their very versatile nutritional capability, adaptability to various hydrocarbon rings, and the possession of the pump mechanism which ejects metal complexes as soon as they enter the cell [47]. In addition, Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli were sensitive to all the complexes, and an inhibitory zone was obtained in the range of 9.0–23.0 mm (Figure 7).In all cases, the metal complexes were found to be more active than the H2L ligand expectedly due to chelation, which reduced the polarity of the metal atom, mainly because of partial sharing of its positive charge with donor groups of the ligand and possible π-electron delocalization on the aromatic rings. This increased the lipophilic character, favouring its permeation into the bacterial membrane, causing the death of the organisms [48].Higher activity of the metal complex was probably due to greater lipophilic nature of the complex. It increased the activity of the metal complex and can be explained on the basis of Overtone’s concept and Tweedy’s chelation theory [49–52]. According to Overtone’s concept of cell permeability, the lipid membrane that surrounds the cell favours the passage of only lipid-soluble materials due to which liposolubility was considered to be an important factor that controls the antimicrobial activity. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups [53, 54]. Furthermore, it increased the delocalization of the π electrons over the whole chelate ring and enhanced the lipophilicity of the complex. This increased lipophilicity enhanced the penetration of the complexes into the lipid membrane and thus blocked the metal-binding sites on enzymes of microorganisms [55]. These metal complexes also disturb the respiration process of the cell and thus block the synthesis of proteins, which restricts further growth of the organism [56].A look at the antibiotic, amikacin, activities (6.0–9.0 mm) against various bacterial isolates relative to the metal complexes (9.0–23.0 mm) showed that the activities of the former were much lower, with optimum activity being about half of metal complexes, against all the bacterial organisms.When the antimicrobial activity of metal complexes is investigated, the following principal factors [57] should be considered: (i) the chelate eﬀect of the ligands; (ii) the nature of the N-donor ligands; (iii) the total charge of the complex; (iv) the existence and the nature of the ion neutralizing the ionic complex; and (v) the nuclearity of the metal center in the complex. This is probably one of the reasons for the diverse antibacterial activity shown by the complexes, while the nature of the metal ion coordinated to the H2L ligand may have a signiﬁcant role to this diversity. In general, all the complexes exhibited better inhibition than free H2L against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli (Table 4). More speciﬁcally, Cd(II) and Co(II) complexes showed the best inhibition among all the complexes in this study, and they were one to twenty-three times more active than the H2L diazo ligand against all the microorganisms used, indicating that the coordination of the H2L ligand to Cd(II) and Co(II) ions enhanced its antimicrobial activity (Figure 7). On the contrary, the rest complexes present higher antimicrobial activity to H2L against the four microorganisms.


### 3.10. Molecular Docking

Targeting the minor groove of DNA through binding to a small molecule has long been considered an important tool in molecular recognition of a speciﬁc DNA sequence. DNA is considered a major biological target for metal complexes that have pharmacological activity. In this context, molecular docking between H2L and [Cd(H2L)Cl2].H2O with four possible biological targets, receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and the crystal structure of Escherichia coli (PDB ID: 3T88), was performed. The structures of these macromolecules were obtained from the Protein Data Bank. The calculated binding energies for the interaction with these proteins are listed in Tables 5 and 6 . The 3D plot curves of docking with free H2L and [Cd(H2L)Cl2].H2O are shown in Figures 8 and 9.Table 5
Energy values obtained in docking calculations of H2L with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88).
ReceptorLigand moietyReceptorInteractionDistanceE (kcal/mol)5JPEN 41O HIS 413 (A)H-donor3.07−2.5N 14NE2 HIS 290 (A)H-acceptor3.33−2.3O 34NH2 ARG 386 (A)H-acceptor3.01−2.8O 36NH2 ARG 261 (A)H-acceptor2.93−1.4O 36NE2 HIS 290 (A)H-acceptor2.99−1.13HB5N 41OD1 ASN 152(X)H-donor2.83−3.1N 42O GLY 9(X)H-donor3.44−0.7O 36OG SER 142(X)H-acceptor2.87−1.56-ringCD2 PHE 192(X)Pi-H3.75−0.63Q8UO 38OE1 GLU 95 (A)H-donor2.90−6.0N 42OE2 GLU 51 (A)H-donor3.00−2.0N 14N THR 91 (A)H-acceptor3.54−1.6O 34NZ LYS 9 (A)H-acceptor3.58−0.6O 34NH2 ARG 102 (A)H-acceptor3.04−2.9O 38OE1 GLU 95 (A)H-donor2.90−6.03T88O 38O PHE 162 (A)H-donor3.33−0.6N 41OD2 ASP 163 (A)H-donor3.11−2.2N 14OH TYR 97 (A)H-acceptor3.18−1.2O 36OH TYR 97 (A)H-acceptor3.38−0.6Table 6
Energy values obtained in docking calculations of [Cd(H2L)Cl2].H2O with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88).
ReceptorLigand moietyReceptorInteractionDistanceE (kcal/mol)5JPEN 42OD2 ASP 373 (A)H-donor2.89−3.1O 48O ARG 386 (A)H-donor3.07−0.7O 38NH2 ARG 386 (A)H-acceptor3.16−2.46-ringCE2 TYR 437 (A)Pi-H4.62−0.63HB5O 48O CYS 185 (X)H-donor2.53−0.2O 48OG SER 142 (X)H-acceptor2.68−4.76-ringCG1 VAL 225 (X)Pi-H4.14−0.73Q8UN 10OE1 GLU 51 (A)H-donor2.72−18N 10OE2 GLU 51 (A)H-donor3.38−0.7CL 36O HIS 48 (A)H-donor3.68−0.7O 48O GLY 116 (A)H-donor2.79−2.2O 50NH1 ARG 125 (A)H-acceptor2.85−4.4N 10OE1 GLU 51 (A)Ionic2.72−6.7N 10OE2 GLU 51 (A)Ionic3.38−2.4N 10OD1 ASP 118 (A)Ionic2.69−7.03T88N 10O PHE 162 (A)H-donor2.89−5.2O 38OH TYR 97 (A)H-acceptor2.68−2.0Figure 8
3D plot of the interaction between the H2L ligand with receptors of (a) 5JPE, (b) 3HB5, (c) 3Q8U, and (d) 3T88.
(a)(b)(c)(d)Figure 9
3D plot of the interaction between Cd-H2L with receptors of (a) 5JPE, (b) 3HB5, (c) 3Q8U, and (d) 3T88.
(a)(b)(c)(d)From these data, the following can be concluded [58–60]:(1)
First, H2L and its Cd(II) complex showed strong binding affinities for all proteins(2)
The [Cd(H2L)Cl2].H2O complex binds stronger than the free ligand to different proteins, see Figure 10(3)
The binding energy for the interaction of the cadmium (II) complex with the receptor protein 3Q8U was three times less than the free diazo ligand confirming the high activity of this compound [58](4)
A closer look at the interactions indicated that the presence of heteroatoms such as oxygen and nitrogen was critical to how the compound binds [59](5)
Generally speaking, H-acceptor interaction dominates the binding of the synthesized complexes to different receptors besides the metal binding to the protein backboneFigure 10
The binding energy of the interaction between H2L and Cd-H2L and receptors of 5JPE, 3HB5, 3Q8U, and 3T88.


## 3.1. Characterization of the Free Diazo Ligand (H2L)

The synthesized diazo ligand (H2L) was soluble in DMF and DMSO solvents and of high melting point which indicated its stability. From the elemental analysis data of the ligand (Supplementary Table 1), it was found that the theoretical values were in agreement with the found values. The IR spectrum of the free diazo ligand was carried out in the range of 4000–400 cm−1, and the most effective bands are listed in Supplementary Table 2. The broad band observed at 3433 cm−1 can be attributed to the ν(O-H) stretching vibration. Also, band observed at 1660 cm−1 and sharp band at 1603 cm−1 can be attributed to C=O stretching vibration [13] and N=N [14], respectively. Also, the ν(COO)asym and ν(COO)sym stretching vibration bands were observed at 1543 and 1384 cm−1, respectively.1H NMR spectrum of the free diazo ligand (H2L) in DMSO-d6 exhibited the following signals: 7.25–7.91 ppm (m, 10H, Ar-H) [13], 4.78 ppm (s, 4H, NH2) [15–17], and 10.12 ppm (s, 2H, COOH). These signals indicated that three different types of protons were present in the diazo ligand. The position of NH2 was confirmed from carrying out the spectra in the deuterated solvent.The mass spectrum of the ligand (H2L) showed an accurate parent molecular ion peak at m/z 404 amu, matched with the theoretical molecular weight. Also, the mass spectrum displayed multiple peaks corresponding to successive degradation of the ligand. The base peak appeared at m/z = 134 due to the [C7H6NO2]+ ion. The other fragments gave the peaks at 104, 132, 136, 164, 240, and 268 amu with various intensities.


## 3.2. Characterization of Metal Complexes

### 3.2.1. Elemental Analyses

The stoichiometry and formulation of the free diazo ligand (H2L) and its metal complexes were confirmed by their elemental analysis (Supplementary Table 1). The metal : ligand ratio was found to be 1 : 1 in all complexes, which has been arrived by estimating the carbon, hydrogen, nitrogen, chloride, and metal contents of the complexes. The elemental analyses of the ligand and its complexes revealed good agreement with the proposed structures. The ligand and its metal complexes have high melting points, and they were found to be air-stable. The diazo ligand was soluble in common organic solvents, and all the complexes were freely soluble in DMF and DMSO but insoluble in methanol, ethanol, and water. The results of elemental analyses suggested formulae [M(H2L)Cl2].nH2O (M = Co(II), Ni(II), Cu(II), Mn(II), Zn(II) and Cd(II)) and [M(H2L)Cl2]Cl.nH2O (M = Cr(III) and Fe(III)).


### 3.2.2. Molar Conductance

The results given in Supplementary Table1 showed that the Cr(III) and Fe(III) complexes had molar conductivity values of 59 and 77 Ω−1 mol−1 cm2, respectively. It is obvious from these data that these chelates were ionic in nature, and they were electrolytes [18], but the molar conductivity values of Mn(II), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) chelates fall in the range of 11–31 Ω−1 mol−1 cm2 indicative of their nonelectrolytic nature [14].


### 3.2.3. IR Spectral Studies

The infrared spectra of the H2L ligand and its metal complexes taken in the region 4000–400 cm−1 are listed in Supplementary Table 2 which gave the most useful assignments for bands diagnostic to the mode of coordination of the ligand. The band observed at 1603 cm−1 in the free diazo ligand (H2L) was assigned to the ν(N=N) stretching vibration. This band was found in the spectra of the complexes at 1610–1628 cm−1. Also, another band observed at 1660 cm−1 assigned to carbonyl of carboxylate ν(C=O) at the free diazo ligand (H2L) was found in the spectra of the complexes at 1649–1656 cm−1 (Supplementary Table 2). The shift in the azo groups of the diazo ligand confirmed the participation of the azo group in chelation.The broad band at 3433 cm−1 observed in the diazo ligand was shifted to 3423–3450 cm−1, which was attributed to the ν(O-H) stretching vibration. This shift can be attributed to the bond formation with metal ions and also the presence of hydrated water molecules. The broadening of this band may be attributed to the presence of the intramolecular hydrogen bond [19].Theνasym(COO−) and νsym(COO−) stretching vibrations were observed at 1543 and 1384 cm −1, respectively, for the free diazo ligand (H2L) [19–26]. Coordination with metal ions via the carboxylate O atom and nitrogen atom was indicated by the shift in the position of the νasym(COO) and νsym(COO) stretching vibration bands to 1534–1555 and1384–1402 cm−1, respectively, for ligand-metal complexes.New bands of low intensity were observed in the far-IR region in the range of 530–578 and 427–445 cm−1 which can be assigned to ν(M-O) and ν(M-N) stretching vibrations, respectively [20–28]. Therefore, from the IR spectral studies, it was concluded that the H2L ligand behaved as a neutral tetradentate ligand with N coordination sites and protonated carboxylate oxygen upon coordinated to the metal ions [20–27].


### 3.2.4.1H NMR Spectral Studies

The1HNMR spectra of the diazo ligand and its Zn(II) complex were recorded in DMSO-d6 solution using tetramethylsilane (TMS) as the internal standard. 1H NMR spectrum of the ligand revealed the following signals at δ = 4.78 (m, 4H, NH2), 7.05–7.94 (m, 10H, Ar-H), and 10.18 ppm (s, 2H, COOH). The shift in the carboxylic acid protons (10.40 ppm) indicated that oxygen of (COOH) groups was coordinated to the Zn(II) ion. However, the singlet signal of the amino group that appeared at δ 4.78 ppm (s, 4H, NH2) means that the two singlet signals assigned to the two NH2 protons remained in their position as in the free diazo ligand (H2L), confirming the remaining of these groups inert towards coordination. The upfield shift in the spectrum of the complex relative to the free ligand can be attributed to the effect of metal ions in chelation.


### 3.2.5. Mass Spectroscopy

The electronic mass spectrum of the [Zn(H2L)Cl2].H2O complex confirmed formation of this complex by showing a molecular ion peak at m/z = 540.81 g/mol corresponding to (M-H2O), which was very close to the expected molecular weight which was 540.50 g/mol, and it also showed a peak at m/z = 403 amu, respectively, corresponding to the ligand moiety. The intensities of these peaks showed the stability of the complex.


### 3.2.6. Spectrophotometric Studies

UV-visible spectral data of the ligand and its chelate (10−4 M) solutions in DMF were recorded from 200 to 700 nm at room temperature using the same solvent as blank. The free ligand (H2L) exhibited broad bands in the UV-visible region at 375 and 271 nm. The first band was assigned to the intraligand n ⟶ π∗ electronic transition. The second broad band around 271 nm was assigned to the π ⟶ π∗ electronic transition. The n ⟶ π∗ and π ⟶ π∗ transition bands were shifted to 339–395 and 261–285 nm, respectively, indicating that the ligand coordinated to metal ions [29, 30].


### 3.2.7. Magnetic Susceptibility and Electronic Spectral Studies

A comparison of the electronic spectra of the free diazo ligand (H2L) coordinated with those of the corresponding metal complexes was discussed. This can be considered as evidence for the complex formation. Additionally, the diffused reflectance spectra of metal complexes showed different bands at different wavelengths; each one was corresponding to a certain transition which suggested the geometry of the complexes. The diffused reflectance spectra of the complexes were dominated by intense intraligand charge-transfer bands.The diffused reﬂectance spectrum of the Cr(III) complex exhibited three bands at 28,489, 24,780, and 13,449 cm−1which may be assigned to the 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P) spin-allowed d–d transitions, respectively. The magnetic moment value was found to be 4.81 BM which indicated the presence of the Cr(III) complex in the octahedral geometry [31, 32].The diffused reﬂectance spectrum of the Mn(II) complex showed three bands at 26,645, 19,490, and 15.660 cm−1 assignable to 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g(D) ⟶ 6A1g transitions, respectively [31]. The magnetic moment value was found to be 5.54 BM which indicated the presence of the Mn(II) complex in the octahedral structure.From the diffused reﬂectance spectrum, it is observed that the Fe(III) chelate exhibited bands at 21,490 and 20,890 cm−1, which may be assigned to the 6A1g ⟶ T2g(G) and 6A1g ⟶ 5T1g transitions in the octahedral geometry of the complex [33]. The observed magnetic moment value of the Fe(III) complex was found to be 5.30 BM. The reflectance spectrum of the Co(II) complex displayed bands near 23,110, 15,240, and 12.845 cm−1, assigned to 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions, respectively. These bands suggested an octahedral geometry for the Co(II) complex. The magnetic moment value of the Co(II) complex was 5.16 BM which was consistent with the high-spin octahedral geometry.The diffused reflectance spectrum of the Ni(II) complex showed absorption bands at 20,680, 14,757, and 13,120 cm−1 which are attributed to 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions, respectively, and assignable to the octahedral geometry for the Ni(II) complex [34]. The magnetic moment value of the Ni(II) complex was 3.80 BM corresponding to the octahedral complex. The diffused reflectance spectrum of the Cu(II) complex exhibited broad bands in ranges 24,729 and 13,290 cm−1 assigned to 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions, respectively. These bands suggested distorted octahedral geometry for the Cu(II) complex. The magnetic moment value was found to be 1.78 BM indicating octahedral geometry [35].The Zn(II) and Cd(II) complexes were diamagnetic. According to the empirical formulae, an octahedral geometry was proposed for these chelates.


### 3.2.8. ESR Studies of the Cu(II)-H2L Complex

The ESR spectrum of the Cu(II) complex was recorded in DMSO at 300 and 77 K (Figure1). The observed order for the Cu(II) complex (A║ (197) > A┴ (98.88); g║ (2.445) > g┴ (2.333)) indicated that the complex exerted an octahedral geometry [35–37]. The observed value of G for the Cu(II) complex (G = 1.421) implied that the exchange coupling was not present, and misalignment was appreciable. The trend g|| > g⊥ > ge (2.0023) showed that the unpaired electron was localized in the dx2-y2 orbital of the Cu(II) ion in the complex [35–38]. The giso (2.0023) value was less than 2.3 which indicated the covalent character of the metal-ligand bond, and the α2 value (0.76) suggested the presence of in-plane covalency. The calculated value (g║/A║) 124 cm for the complex was consistent with the slightly distorted structure. The orbital reduction factors K║ and K⊥ were estimated from the following expressions: K║ = (g║ −2.0023) ΔE/8λ and K⊥ = (g║ −2.0023) ΔE/2λ, λ = −828 cm−1 (spin-orbit coupling constant for the free ion), and K║ (0.718) > K⊥ (0.45) showed poor in-plane π bonding.Figure 1
ESR spectrum of the [Cu(H2L)Cl2].3/2H2O complex


### 3.2.9. Powder X-Ray Diffraction and SEM

The XRD patterns indicated crystalline nature for the H2L diazo ligand and the Cd(II) complex only. It can be easily seen that the pattern of the H2L diazo ligand differed from its metal complexes, which may be attributed to the formation of a well-deﬁned distorted crystalline structure. Probably, this behavior was due to the incorporation of water molecules into the coordination sphere. On comparing the XRD spectra of the chelates with the XRD spectra of the free ligand, it was concluded that all chelates under study can be considered to have amorphous structures as they lack sharp peaks except the Cd(II) complex where it had a crystalline structure. Therefore, the nonsimilarity of the XRD pattern between the metal ions and chelates suggested that these chelates had different phase structures than the free diazo ligand (H2L). Such facts suggested that H2L and the Cd(II) complex were crystalline, while its Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), and Zn(II) complexes were amorphous.The average crystallite size (ξ) can be calculated from the XRD pattern according to the Debye–Scherrer equation [39, 40]:(1)ξ=Kλβ1/2cosθ.The equation uses the reference peak width at angleθ, where λ is the wavelength of the X-ray radiation (1.541874 Å), K is a constant taken as 0.95 for organic compounds [37], and β1/2 is the width at half maximum of the reference diffraction peak measured in radians. The dislocation density, δ, is the number of dislocation lines per unit area of the crystal. The value of δ is related to the average particle diameter (ξ) by the following relation [41, 42]:(2)δ=1ξ2.The value ofδ was calculated and found to be 33.78 and 26.77 nm, and the value of δ was 8.76 × 10−4 and 1.39 × 10−3 nm−2 for H2L and Cd(II) complex, respectively, giving evidence for the formation of the free azo ligand and its cadmium(II) complex in a nanosize.The SEM micrographs of H2L and its Cd(II) are given in Figure 2. There is a significant difference in the morphological structures, and the particle size is in the nanometer range. The particle size of the ligand was 64–67 nm, while the Cd(II) complex was 20–32 nm.Figure 2
The SEM micrographs of (a) Cd(II)-H2L and (b) H2L.


### 3.2.10. Contact Angle

For two centuries or more of research, wetting phenomena still remain popular topics of investigations. Young’s static and dynamic contact angle equations [6, 7] are the most common and relevant parameters describing the wetting property of a surface with respect to liquid water:(3)cosθY=σsv−σs1σ1v,where θY = Young’s contact angle, σsv = surface tension (energy per unit surface) of the solid-vapor interface, σsl = surface tension of the solid-liquid interface, and σlv = surface tension of the liquid-vapor interface.The mean contact angles were measured for H2L and its metal complexes (Supplementary Figure 3). The data listed in Table 1 proved that the diazo ligand and all the complexes are hydrophobic, and the Cd(II) complex can be considered as a superhydrophobic material. Superhydrophobic surfaces are actively studied across a wide range of applications and industries and are now finding increased use in the biomedical areas. A partial list includes clothing that will be both breathable and water repellant, umbrellas that stayed completely dry, building materials, paints, epoxies, goods transport, and anticorrosion. Superhydrophobic materials such as this promising nano-Cd(II) complex exhibit a number of unique properties that arise from the high roughness of a low surface energy material that stabilizes a nonwetted state [6, 7].Table 1
The mean contact angle of H2L and its complexes.
CompoundMean contact angle (°)Wettability classificationSurface morphology[Cr(H2L)Cl2]Cl.2H2O137HydrophobicRough and porous[Mn(H2L)Cl2]128HydrophobicRough and porous[Fe(H2L)Cl2]Cl.H2O142HydrophobicRough and porous[Co(H2L)Cl2]123HydrophobicRough and porous[Ni(H2L)Cl2]132HydrophobicRough and porous[Cu(H2L)Cl2].3/2H2O133HydrophobicRough and porous[Zn(H2L)Cl2].H2O128HydrophobicRough and porous[Cd(H2L)Cl2].H2O153SuperhydrophobicVery rough and very porousH2L azo dye ligand103HydrophobicRough and porous


### 3.2.11. Thermal Analyses

Thermogravimetric analyses of the metal complexes under investigation were used to get information about the thermal stability of these new complexes, to decide whether the water molecules (if present) were inside or outside the inner coordination sphere of the central metal ion, and to suggest a general scheme for thermal decomposition of these chelates. In the present investigation, heating rates were suitably controlled at 10°C min−1 under nitrogen atmosphere, and the weight loss was measured from the ambient temperature up to ∼1000°C. The TG curves of the complexes are listed in Table 2. The weight losses for each chelate were calculated within the corresponding temperature ranges.Table 2
Thermoanalytical results (TG and DTG) of the H2L ligand and its metal complexes.
ComplexTG range (oC)DTG maxn∗Mass loss estim (calcd)% (total mass loss)AssignmentResiduesDiazo ligand (H2L)30–355216136.53 (38.61)Loss of 4NH3 and 2CO2—355–1000476,602261.76 (61.38) 100.0 (99.90)Loss of C18H14N2[Cr(H2L)Cl2]Cl.2H2O45–1157918.88 (8.53)Loss of 2H2O and NH31/2 Cr2O3115–310210130.00 (29.90)Loss of 3/2Cl2, 2N2, and CH4310–990377,865241.30 (43.10) 80.18 (81.53)Loss of C17H12NO2.5[Mn(H2L)Cl2]30–435234158.98 (58.86)Loss of 2N2,Cl2, and C9H5N27C + MnO435–1000592120.88 (20.18) 79.66 (79.04)Loss of C4H11O3[Fe(H2L)Cl2]Cl.H2O45–1007513.29 (3.07)Loss of H2O5C + 1/2Fe2O3100–355244147.89 (48.33)Loss of 3/2Cl2, 3N2, 3CH4, and CO2355–1000485,565,803325.37 (24.63) 76.62 (76.03)Loss of C11H4O0.5[Co(H2L)Cl2]30–260177112.23 (12.45)Loss of 2NH3 and1/2Cl22C + CoO260–47536416.33 (6.64)Loss of 1/2 Cl2475–890640160.3 (60.03) 78.91 (79.09)Loss of C18H13N5O[Ni(H2L)Cl2]40–290199131.24 (31.27)Loss of 4NH3,N2, and Cl28C + NiO290–970395,883235.57 (36.70) 67.85 (67.95)Loss of C12H4O3[Cu(H2L)Cl2].3/2H2O45–1458814.71 (4.77)Loss of 3/2H2O2C + CuO145–990287177.67 (76.92) 82.38 (81.70)Loss of C18H16Cl2N6O3[Zn(H2L)Cl2].H2O45–1308615.42 (5.72)Loss of H2O and 1/2N23C + ZnO130–270205110.36 (10.26)Loss of 2N2270–1000348,467263.4 (63.20) 79.17 (79.18)Loss of C17H16Cl2NO3[Cd(H2L)Cl2].H2O40–1309812.95 (2.90)Loss of H2OCdO130–355222125.63 (25.62)Loss of 3N2 and Cl2355–1000414,591,833350.15 (50.25) 78.73 (78.77)Loss of C20H16O3n∗ = number of decomposition steps.The thermal decomposition process of H2L involved three decomposition steps. Decomposition of H2L started at 30°C and finished at 100°C with three stages. The first stage of decomposition involved the removal of four molecules of NH3 and two molecules of carbon dioxide gas in the 30–355°C temperature range and was accompanied by a weight loss of 36.53% (calcd. = 36.41%). The second and third stages of decomposition occurred in the 355–1000°C temperature range, corresponding to the loss of the C18H4N2 molecule and were accompanied by a weight loss of 62.76% (calcd. = 62.59%).Decomposition of the [Cr(H2L)Cl2]Cl.2H2O complex started at 45°C and ﬁnished at 1000°C with four stages at four maxima at 79, 210, 377, and 865°C and was accompanied by a weight loss of 80.29%, corresponding to the loss of 2H2O, 3/2Cl2, 2N2, CH4, and C17H12NO2.5 molecules which was very close to the calculated value of 81.53%.Thermogravimetric curve for the [Mn(H2L)Cl2] complex showed two weight loss events. The ﬁrst step of decomposition occurred within the range of 30–435°C, with a maximum temperature at 234°C, and it corresponds to the loss of Cl2, 2N2 gases and C9H5N2 molecules. The second step of decomposition occurred in the range of 435–1000°C with the maximum temperature at 595oC. There were signiﬁcant mass loss events due to the decomposition of the C4H11O3molecule (Table 2).The thermal decomposition of the [Fe(H2L)Cl2]Cl.H2O complex proceeded via five degradation steps. The ﬁrst step of decomposition occurred within the temperature range from 45 to 100°C, with a maximum temperature at 75°C, and it corresponds to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 100–355°C, with one maxima at 244°C, and corresponds to the elimination of 3/2Cl2, 3N2, 3CH4, and CO2 with a mass loss of 47.89% (calcd. = 48.33%). The last three steps of decomposition occurred within the temperature range of 355–1000°C with three maxima at 485, 565, and 803°C and were simultaneously decomposed to ferric oxide contaminated with carbon as a ﬁnal product. They were accompanied by a weight loss of 25.37% (calcd. = 24.63%) and corresponded to the loss of the C11H4O0.5 molecule. The actual weight loss from these five steps was 76.62% which was close to the calculated value of 76.03%.The thermal degradation of the [Co(H2L)Cl2] complex took place in three degradation stages within the temperature range from 30 to 990oC. The ﬁrst stage of decomposition occurred at maximum 177°C and was accompanied by a weight loss of 12.23% (calcd. = 12.45%) corresponding to the loss of 2NH3 and 1/2Cl2 gases. The second and third steps of decomposition occurred at two maxima 364 and 640°C and were accompanied by a weight loss of 6.33% (calcd. = 6.64%) and 60.30% (calcd. = 60.03%), respectively, corresponding to the loss of 1/2Cl2 gas and C18H3N5O molecule giving CoO contaminated with carbon as a ﬁnal product. Theoretically, the weight loss in these steps was 78.91% which agrees with the experimental value of 79.09%.The [Ni(H2L)Cl2] complex lost upon heating 4NH3, Cl2, and N2 gases in the ﬁrst step of decomposition within the temperature range of 40–290oC. The last two steps of decomposition occurred at two maxima at 395 and 880°C. These steps were associated with the loss of the C12H4O3 molecule forming NiO contaminated with carbon as a ﬁnal product. The actual total weight loss of 67.85% was in agreement with the calculated total weight loss value of 67.95%.The complex [Cu(H2L)Cl2].3/2H2O lost upon heating all water molecules at one maxima 90oC with an estimated weight loss of 4.70% (calcd. = 4.77%). The dehydrated Cu(II) complex was simultaneously decomposed to CuO and 2C at the 287oC maximum temperature. The total weight loss amounted to 82.38% (calcd. = 81.74%).The [Zn(H2L)Cl2].H2O complex was thermally decomposed into four steps within the temperature range from 45 to 1000oC. The ﬁrst decomposition step with an estimated mass loss of 5.42% (calcd. = 5.72%) occurred within the temperature range from 45 to 130oC. This step may be attributed to the liberation of the hydrated water molecules and NH3 gas. The remaining three decomposition steps of the Zn(II) complex were observed within the temperature range from 130 to 1000oC. Estimated mass losses of 73.76% (calcd. = 74.42%) were reasonably accounted for the removal of two 2N2 gases and a C17H16Cl2NO3 molecule in two steps.TG curve of the [Cd(H2L)Cl2].H2O complex displayed a five-step decomposition. The ﬁrst step of decomposition occurred within the temperature range from 40 to 130oC, with a maximum temperature at 98 oC, and corresponded to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 130–355 oC, with one maxima at 222oC, and corresponded to the elimination of Cl2 and 3N2 gases with a mass loss of 25.62% (calcd. = 25.63%). The last three steps of decomposition occurred within the temperature range of 355–1000oC with three maxima at 414, 591, and 833oC and were simultaneously decomposed to CdO as a ﬁnal product. They were accompanied by a weight loss of 50.25% (calcd. = 50.15%) corresponding to the loss of the C20H16O3 molecule. The weight loss from these five steps was 78.77% which was close to the calculated value of 78.73%.


## 3.2.1. Elemental Analyses

The stoichiometry and formulation of the free diazo ligand (H2L) and its metal complexes were confirmed by their elemental analysis (Supplementary Table 1). The metal : ligand ratio was found to be 1 : 1 in all complexes, which has been arrived by estimating the carbon, hydrogen, nitrogen, chloride, and metal contents of the complexes. The elemental analyses of the ligand and its complexes revealed good agreement with the proposed structures. The ligand and its metal complexes have high melting points, and they were found to be air-stable. The diazo ligand was soluble in common organic solvents, and all the complexes were freely soluble in DMF and DMSO but insoluble in methanol, ethanol, and water. The results of elemental analyses suggested formulae [M(H2L)Cl2].nH2O (M = Co(II), Ni(II), Cu(II), Mn(II), Zn(II) and Cd(II)) and [M(H2L)Cl2]Cl.nH2O (M = Cr(III) and Fe(III)).


## 3.2.2. Molar Conductance

The results given in Supplementary Table1 showed that the Cr(III) and Fe(III) complexes had molar conductivity values of 59 and 77 Ω−1 mol−1 cm2, respectively. It is obvious from these data that these chelates were ionic in nature, and they were electrolytes [18], but the molar conductivity values of Mn(II), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) chelates fall in the range of 11–31 Ω−1 mol−1 cm2 indicative of their nonelectrolytic nature [14].


## 3.2.3. IR Spectral Studies

The infrared spectra of the H2L ligand and its metal complexes taken in the region 4000–400 cm−1 are listed in Supplementary Table 2 which gave the most useful assignments for bands diagnostic to the mode of coordination of the ligand. The band observed at 1603 cm−1 in the free diazo ligand (H2L) was assigned to the ν(N=N) stretching vibration. This band was found in the spectra of the complexes at 1610–1628 cm−1. Also, another band observed at 1660 cm−1 assigned to carbonyl of carboxylate ν(C=O) at the free diazo ligand (H2L) was found in the spectra of the complexes at 1649–1656 cm−1 (Supplementary Table 2). The shift in the azo groups of the diazo ligand confirmed the participation of the azo group in chelation.The broad band at 3433 cm−1 observed in the diazo ligand was shifted to 3423–3450 cm−1, which was attributed to the ν(O-H) stretching vibration. This shift can be attributed to the bond formation with metal ions and also the presence of hydrated water molecules. The broadening of this band may be attributed to the presence of the intramolecular hydrogen bond [19].Theνasym(COO−) and νsym(COO−) stretching vibrations were observed at 1543 and 1384 cm −1, respectively, for the free diazo ligand (H2L) [19–26]. Coordination with metal ions via the carboxylate O atom and nitrogen atom was indicated by the shift in the position of the νasym(COO) and νsym(COO) stretching vibration bands to 1534–1555 and1384–1402 cm−1, respectively, for ligand-metal complexes.New bands of low intensity were observed in the far-IR region in the range of 530–578 and 427–445 cm−1 which can be assigned to ν(M-O) and ν(M-N) stretching vibrations, respectively [20–28]. Therefore, from the IR spectral studies, it was concluded that the H2L ligand behaved as a neutral tetradentate ligand with N coordination sites and protonated carboxylate oxygen upon coordinated to the metal ions [20–27].


## 3.2.4.1H NMR Spectral Studies

The1HNMR spectra of the diazo ligand and its Zn(II) complex were recorded in DMSO-d6 solution using tetramethylsilane (TMS) as the internal standard. 1H NMR spectrum of the ligand revealed the following signals at δ = 4.78 (m, 4H, NH2), 7.05–7.94 (m, 10H, Ar-H), and 10.18 ppm (s, 2H, COOH). The shift in the carboxylic acid protons (10.40 ppm) indicated that oxygen of (COOH) groups was coordinated to the Zn(II) ion. However, the singlet signal of the amino group that appeared at δ 4.78 ppm (s, 4H, NH2) means that the two singlet signals assigned to the two NH2 protons remained in their position as in the free diazo ligand (H2L), confirming the remaining of these groups inert towards coordination. The upfield shift in the spectrum of the complex relative to the free ligand can be attributed to the effect of metal ions in chelation.


## 3.2.5. Mass Spectroscopy

The electronic mass spectrum of the [Zn(H2L)Cl2].H2O complex confirmed formation of this complex by showing a molecular ion peak at m/z = 540.81 g/mol corresponding to (M-H2O), which was very close to the expected molecular weight which was 540.50 g/mol, and it also showed a peak at m/z = 403 amu, respectively, corresponding to the ligand moiety. The intensities of these peaks showed the stability of the complex.


## 3.2.6. Spectrophotometric Studies

UV-visible spectral data of the ligand and its chelate (10−4 M) solutions in DMF were recorded from 200 to 700 nm at room temperature using the same solvent as blank. The free ligand (H2L) exhibited broad bands in the UV-visible region at 375 and 271 nm. The first band was assigned to the intraligand n ⟶ π∗ electronic transition. The second broad band around 271 nm was assigned to the π ⟶ π∗ electronic transition. The n ⟶ π∗ and π ⟶ π∗ transition bands were shifted to 339–395 and 261–285 nm, respectively, indicating that the ligand coordinated to metal ions [29, 30].


## 3.2.7. Magnetic Susceptibility and Electronic Spectral Studies

A comparison of the electronic spectra of the free diazo ligand (H2L) coordinated with those of the corresponding metal complexes was discussed. This can be considered as evidence for the complex formation. Additionally, the diffused reflectance spectra of metal complexes showed different bands at different wavelengths; each one was corresponding to a certain transition which suggested the geometry of the complexes. The diffused reflectance spectra of the complexes were dominated by intense intraligand charge-transfer bands.The diffused reﬂectance spectrum of the Cr(III) complex exhibited three bands at 28,489, 24,780, and 13,449 cm−1which may be assigned to the 4A2g(F) ⟶ 4T2g(F), 4A2g(F) ⟶ 4T1g(F), and 4A2g(F) ⟶ 4T2g(P) spin-allowed d–d transitions, respectively. The magnetic moment value was found to be 4.81 BM which indicated the presence of the Cr(III) complex in the octahedral geometry [31, 32].The diffused reﬂectance spectrum of the Mn(II) complex showed three bands at 26,645, 19,490, and 15.660 cm−1 assignable to 4T1g ⟶ 6A1g, 4T2g(G) ⟶ 6A1g, and 4T1g(D) ⟶ 6A1g transitions, respectively [31]. The magnetic moment value was found to be 5.54 BM which indicated the presence of the Mn(II) complex in the octahedral structure.From the diffused reﬂectance spectrum, it is observed that the Fe(III) chelate exhibited bands at 21,490 and 20,890 cm−1, which may be assigned to the 6A1g ⟶ T2g(G) and 6A1g ⟶ 5T1g transitions in the octahedral geometry of the complex [33]. The observed magnetic moment value of the Fe(III) complex was found to be 5.30 BM. The reflectance spectrum of the Co(II) complex displayed bands near 23,110, 15,240, and 12.845 cm−1, assigned to 4T1g(F) ⟶ 4T2g(F), 4T1g(F) ⟶ 4A2g(F), and 4T1g(F) ⟶ 4T1g(P) transitions, respectively. These bands suggested an octahedral geometry for the Co(II) complex. The magnetic moment value of the Co(II) complex was 5.16 BM which was consistent with the high-spin octahedral geometry.The diffused reflectance spectrum of the Ni(II) complex showed absorption bands at 20,680, 14,757, and 13,120 cm−1 which are attributed to 3A2g(F) ⟶ 3T2g(F), 3A2g(F) ⟶ 3T1g(F), and 3A2g(F) ⟶ 3T1g transitions, respectively, and assignable to the octahedral geometry for the Ni(II) complex [34]. The magnetic moment value of the Ni(II) complex was 3.80 BM corresponding to the octahedral complex. The diffused reflectance spectrum of the Cu(II) complex exhibited broad bands in ranges 24,729 and 13,290 cm−1 assigned to 2B1g ⟶ 2B2g, 2B1g ⟶ 2Eg and 2B1g ⟶ 2A1g transitions, respectively. These bands suggested distorted octahedral geometry for the Cu(II) complex. The magnetic moment value was found to be 1.78 BM indicating octahedral geometry [35].The Zn(II) and Cd(II) complexes were diamagnetic. According to the empirical formulae, an octahedral geometry was proposed for these chelates.


## 3.2.8. ESR Studies of the Cu(II)-H2L Complex

The ESR spectrum of the Cu(II) complex was recorded in DMSO at 300 and 77 K (Figure1). The observed order for the Cu(II) complex (A║ (197) > A┴ (98.88); g║ (2.445) > g┴ (2.333)) indicated that the complex exerted an octahedral geometry [35–37]. The observed value of G for the Cu(II) complex (G = 1.421) implied that the exchange coupling was not present, and misalignment was appreciable. The trend g|| > g⊥ > ge (2.0023) showed that the unpaired electron was localized in the dx2-y2 orbital of the Cu(II) ion in the complex [35–38]. The giso (2.0023) value was less than 2.3 which indicated the covalent character of the metal-ligand bond, and the α2 value (0.76) suggested the presence of in-plane covalency. The calculated value (g║/A║) 124 cm for the complex was consistent with the slightly distorted structure. The orbital reduction factors K║ and K⊥ were estimated from the following expressions: K║ = (g║ −2.0023) ΔE/8λ and K⊥ = (g║ −2.0023) ΔE/2λ, λ = −828 cm−1 (spin-orbit coupling constant for the free ion), and K║ (0.718) > K⊥ (0.45) showed poor in-plane π bonding.Figure 1
ESR spectrum of the [Cu(H2L)Cl2].3/2H2O complex


## 3.2.9. Powder X-Ray Diffraction and SEM

The XRD patterns indicated crystalline nature for the H2L diazo ligand and the Cd(II) complex only. It can be easily seen that the pattern of the H2L diazo ligand differed from its metal complexes, which may be attributed to the formation of a well-deﬁned distorted crystalline structure. Probably, this behavior was due to the incorporation of water molecules into the coordination sphere. On comparing the XRD spectra of the chelates with the XRD spectra of the free ligand, it was concluded that all chelates under study can be considered to have amorphous structures as they lack sharp peaks except the Cd(II) complex where it had a crystalline structure. Therefore, the nonsimilarity of the XRD pattern between the metal ions and chelates suggested that these chelates had different phase structures than the free diazo ligand (H2L). Such facts suggested that H2L and the Cd(II) complex were crystalline, while its Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), and Zn(II) complexes were amorphous.The average crystallite size (ξ) can be calculated from the XRD pattern according to the Debye–Scherrer equation [39, 40]:(1)ξ=Kλβ1/2cosθ.The equation uses the reference peak width at angleθ, where λ is the wavelength of the X-ray radiation (1.541874 Å), K is a constant taken as 0.95 for organic compounds [37], and β1/2 is the width at half maximum of the reference diffraction peak measured in radians. The dislocation density, δ, is the number of dislocation lines per unit area of the crystal. The value of δ is related to the average particle diameter (ξ) by the following relation [41, 42]:(2)δ=1ξ2.The value ofδ was calculated and found to be 33.78 and 26.77 nm, and the value of δ was 8.76 × 10−4 and 1.39 × 10−3 nm−2 for H2L and Cd(II) complex, respectively, giving evidence for the formation of the free azo ligand and its cadmium(II) complex in a nanosize.The SEM micrographs of H2L and its Cd(II) are given in Figure 2. There is a significant difference in the morphological structures, and the particle size is in the nanometer range. The particle size of the ligand was 64–67 nm, while the Cd(II) complex was 20–32 nm.Figure 2
The SEM micrographs of (a) Cd(II)-H2L and (b) H2L.


## 3.2.10. Contact Angle

For two centuries or more of research, wetting phenomena still remain popular topics of investigations. Young’s static and dynamic contact angle equations [6, 7] are the most common and relevant parameters describing the wetting property of a surface with respect to liquid water:(3)cosθY=σsv−σs1σ1v,where θY = Young’s contact angle, σsv = surface tension (energy per unit surface) of the solid-vapor interface, σsl = surface tension of the solid-liquid interface, and σlv = surface tension of the liquid-vapor interface.The mean contact angles were measured for H2L and its metal complexes (Supplementary Figure 3). The data listed in Table 1 proved that the diazo ligand and all the complexes are hydrophobic, and the Cd(II) complex can be considered as a superhydrophobic material. Superhydrophobic surfaces are actively studied across a wide range of applications and industries and are now finding increased use in the biomedical areas. A partial list includes clothing that will be both breathable and water repellant, umbrellas that stayed completely dry, building materials, paints, epoxies, goods transport, and anticorrosion. Superhydrophobic materials such as this promising nano-Cd(II) complex exhibit a number of unique properties that arise from the high roughness of a low surface energy material that stabilizes a nonwetted state [6, 7].Table 1
The mean contact angle of H2L and its complexes.
CompoundMean contact angle (°)Wettability classificationSurface morphology[Cr(H2L)Cl2]Cl.2H2O137HydrophobicRough and porous[Mn(H2L)Cl2]128HydrophobicRough and porous[Fe(H2L)Cl2]Cl.H2O142HydrophobicRough and porous[Co(H2L)Cl2]123HydrophobicRough and porous[Ni(H2L)Cl2]132HydrophobicRough and porous[Cu(H2L)Cl2].3/2H2O133HydrophobicRough and porous[Zn(H2L)Cl2].H2O128HydrophobicRough and porous[Cd(H2L)Cl2].H2O153SuperhydrophobicVery rough and very porousH2L azo dye ligand103HydrophobicRough and porous


## 3.2.11. Thermal Analyses

Thermogravimetric analyses of the metal complexes under investigation were used to get information about the thermal stability of these new complexes, to decide whether the water molecules (if present) were inside or outside the inner coordination sphere of the central metal ion, and to suggest a general scheme for thermal decomposition of these chelates. In the present investigation, heating rates were suitably controlled at 10°C min−1 under nitrogen atmosphere, and the weight loss was measured from the ambient temperature up to ∼1000°C. The TG curves of the complexes are listed in Table 2. The weight losses for each chelate were calculated within the corresponding temperature ranges.Table 2
Thermoanalytical results (TG and DTG) of the H2L ligand and its metal complexes.
ComplexTG range (oC)DTG maxn∗Mass loss estim (calcd)% (total mass loss)AssignmentResiduesDiazo ligand (H2L)30–355216136.53 (38.61)Loss of 4NH3 and 2CO2—355–1000476,602261.76 (61.38) 100.0 (99.90)Loss of C18H14N2[Cr(H2L)Cl2]Cl.2H2O45–1157918.88 (8.53)Loss of 2H2O and NH31/2 Cr2O3115–310210130.00 (29.90)Loss of 3/2Cl2, 2N2, and CH4310–990377,865241.30 (43.10) 80.18 (81.53)Loss of C17H12NO2.5[Mn(H2L)Cl2]30–435234158.98 (58.86)Loss of 2N2,Cl2, and C9H5N27C + MnO435–1000592120.88 (20.18) 79.66 (79.04)Loss of C4H11O3[Fe(H2L)Cl2]Cl.H2O45–1007513.29 (3.07)Loss of H2O5C + 1/2Fe2O3100–355244147.89 (48.33)Loss of 3/2Cl2, 3N2, 3CH4, and CO2355–1000485,565,803325.37 (24.63) 76.62 (76.03)Loss of C11H4O0.5[Co(H2L)Cl2]30–260177112.23 (12.45)Loss of 2NH3 and1/2Cl22C + CoO260–47536416.33 (6.64)Loss of 1/2 Cl2475–890640160.3 (60.03) 78.91 (79.09)Loss of C18H13N5O[Ni(H2L)Cl2]40–290199131.24 (31.27)Loss of 4NH3,N2, and Cl28C + NiO290–970395,883235.57 (36.70) 67.85 (67.95)Loss of C12H4O3[Cu(H2L)Cl2].3/2H2O45–1458814.71 (4.77)Loss of 3/2H2O2C + CuO145–990287177.67 (76.92) 82.38 (81.70)Loss of C18H16Cl2N6O3[Zn(H2L)Cl2].H2O45–1308615.42 (5.72)Loss of H2O and 1/2N23C + ZnO130–270205110.36 (10.26)Loss of 2N2270–1000348,467263.4 (63.20) 79.17 (79.18)Loss of C17H16Cl2NO3[Cd(H2L)Cl2].H2O40–1309812.95 (2.90)Loss of H2OCdO130–355222125.63 (25.62)Loss of 3N2 and Cl2355–1000414,591,833350.15 (50.25) 78.73 (78.77)Loss of C20H16O3n∗ = number of decomposition steps.The thermal decomposition process of H2L involved three decomposition steps. Decomposition of H2L started at 30°C and finished at 100°C with three stages. The first stage of decomposition involved the removal of four molecules of NH3 and two molecules of carbon dioxide gas in the 30–355°C temperature range and was accompanied by a weight loss of 36.53% (calcd. = 36.41%). The second and third stages of decomposition occurred in the 355–1000°C temperature range, corresponding to the loss of the C18H4N2 molecule and were accompanied by a weight loss of 62.76% (calcd. = 62.59%).Decomposition of the [Cr(H2L)Cl2]Cl.2H2O complex started at 45°C and ﬁnished at 1000°C with four stages at four maxima at 79, 210, 377, and 865°C and was accompanied by a weight loss of 80.29%, corresponding to the loss of 2H2O, 3/2Cl2, 2N2, CH4, and C17H12NO2.5 molecules which was very close to the calculated value of 81.53%.Thermogravimetric curve for the [Mn(H2L)Cl2] complex showed two weight loss events. The ﬁrst step of decomposition occurred within the range of 30–435°C, with a maximum temperature at 234°C, and it corresponds to the loss of Cl2, 2N2 gases and C9H5N2 molecules. The second step of decomposition occurred in the range of 435–1000°C with the maximum temperature at 595oC. There were signiﬁcant mass loss events due to the decomposition of the C4H11O3molecule (Table 2).The thermal decomposition of the [Fe(H2L)Cl2]Cl.H2O complex proceeded via five degradation steps. The ﬁrst step of decomposition occurred within the temperature range from 45 to 100°C, with a maximum temperature at 75°C, and it corresponds to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 100–355°C, with one maxima at 244°C, and corresponds to the elimination of 3/2Cl2, 3N2, 3CH4, and CO2 with a mass loss of 47.89% (calcd. = 48.33%). The last three steps of decomposition occurred within the temperature range of 355–1000°C with three maxima at 485, 565, and 803°C and were simultaneously decomposed to ferric oxide contaminated with carbon as a ﬁnal product. They were accompanied by a weight loss of 25.37% (calcd. = 24.63%) and corresponded to the loss of the C11H4O0.5 molecule. The actual weight loss from these five steps was 76.62% which was close to the calculated value of 76.03%.The thermal degradation of the [Co(H2L)Cl2] complex took place in three degradation stages within the temperature range from 30 to 990oC. The ﬁrst stage of decomposition occurred at maximum 177°C and was accompanied by a weight loss of 12.23% (calcd. = 12.45%) corresponding to the loss of 2NH3 and 1/2Cl2 gases. The second and third steps of decomposition occurred at two maxima 364 and 640°C and were accompanied by a weight loss of 6.33% (calcd. = 6.64%) and 60.30% (calcd. = 60.03%), respectively, corresponding to the loss of 1/2Cl2 gas and C18H3N5O molecule giving CoO contaminated with carbon as a ﬁnal product. Theoretically, the weight loss in these steps was 78.91% which agrees with the experimental value of 79.09%.The [Ni(H2L)Cl2] complex lost upon heating 4NH3, Cl2, and N2 gases in the ﬁrst step of decomposition within the temperature range of 40–290oC. The last two steps of decomposition occurred at two maxima at 395 and 880°C. These steps were associated with the loss of the C12H4O3 molecule forming NiO contaminated with carbon as a ﬁnal product. The actual total weight loss of 67.85% was in agreement with the calculated total weight loss value of 67.95%.The complex [Cu(H2L)Cl2].3/2H2O lost upon heating all water molecules at one maxima 90oC with an estimated weight loss of 4.70% (calcd. = 4.77%). The dehydrated Cu(II) complex was simultaneously decomposed to CuO and 2C at the 287oC maximum temperature. The total weight loss amounted to 82.38% (calcd. = 81.74%).The [Zn(H2L)Cl2].H2O complex was thermally decomposed into four steps within the temperature range from 45 to 1000oC. The ﬁrst decomposition step with an estimated mass loss of 5.42% (calcd. = 5.72%) occurred within the temperature range from 45 to 130oC. This step may be attributed to the liberation of the hydrated water molecules and NH3 gas. The remaining three decomposition steps of the Zn(II) complex were observed within the temperature range from 130 to 1000oC. Estimated mass losses of 73.76% (calcd. = 74.42%) were reasonably accounted for the removal of two 2N2 gases and a C17H16Cl2NO3 molecule in two steps.TG curve of the [Cd(H2L)Cl2].H2O complex displayed a five-step decomposition. The ﬁrst step of decomposition occurred within the temperature range from 40 to 130oC, with a maximum temperature at 98 oC, and corresponded to the loss of one water molecule of hydration. The second step of decomposition occurred in the range of 130–355 oC, with one maxima at 222oC, and corresponded to the elimination of Cl2 and 3N2 gases with a mass loss of 25.62% (calcd. = 25.63%). The last three steps of decomposition occurred within the temperature range of 355–1000oC with three maxima at 414, 591, and 833oC and were simultaneously decomposed to CdO as a ﬁnal product. They were accompanied by a weight loss of 50.25% (calcd. = 50.15%) corresponding to the loss of the C20H16O3 molecule. The weight loss from these five steps was 78.77% which was close to the calculated value of 78.73%.


## 3.3. Geometry Optimization

The geometric parameters (bond lengths and bond angles) of the optimized structures of the H2L azo ligand and its Cd(II) complex (Figure 3) are listed in Supplementary Table 5. From the analysis of these data, the following remarks were found:(1)
There was an elongation in the coordination bonds after complexation, and a large variation in N(11)-N(14), N(12)-N(13), O(35)-H(40), and O(38)-H(39) bond lengths of the ligand occurred upon complexation. It became slightly longer confirming that the coordination took place via azo N atoms and protonated O groups.(2)
All the active groups taking part in coordination have bonds longer than those already existed in the ligands (for example, N=N).(3)
The bond angles of the free ligand moiety were altered somewhat upon coordination; also, the bond angles surrounding the central cadmium atom in the complex lie in the range reported for octahedral geometry [43].(4)
The decrease in the metal-chloride angles may be attributed to the intramolecular hydrogen bond [43].Figure 3
The optimized structure of (a) H2L and (b) [Cd(H2L)Cl2].H2O.
(a)(b)


## 3.4. Molecular Electrostatic Potential (MEP)

The MEP is a plot of electrostatic potential mapped onto the constant electron density surface. It is also very useful in the research of the molecular structure with its physiochemical property relationship. It is also can be used to understand the molecular interactions and predict the reactive sites for electrophilic and nucleophilic attacks [43]. The MEP is directly related to chemical reactivity. Red color represented the regions of negative electrostatic potential. The negative regions are related to electrophilic reactivity. However, the electron-poor region has blue color (favor site for the nucleophilic attack) [43, 44], but the region with green color points to the neutral electrostatic potential region. Figure 4 represents 3D plots of the MEP for the parent ligand and its Cd(II) complex. It was found that oxygen and nitrogen atoms of free H2L were surrounded by a greater negative charge surface, making these sites potentially more favorable for the electrophilic attack (Figure 4(a)), where a greater negative charge was surrounded to the metal center (Figure 4(b)) confirming the complexation process.Figure 4
Molecular electrostatic potential map of (a) H2L and (b) [Cd(H2L)Cl2].H2O. The electron density isosurface is 0.004 a.u.
(a)(b)


## 3.5. Mulliken Charges

Supplementary Table6 and Supplementary Figure 6 clarify Mulliken atomic charges for some important atoms. The calculated charge value for N13 and N14 is −0.08 and −0.007 a.u., respectively, in contrast to the expected one (−1 a.u.), which indicated electron transfer from these atoms to the metal center. Furthermore, decrease of oxygen and chlorine formal charge from the expected values is a reason for migration of electrons to the cadmium center in order to forming the complex. Finally, this charge transfer from donor atoms to the metal center causes the calculated value to decrease to −0.756 a.u. against the formal charge of +2 for the cadmium atom [43, 44].


## 3.6. Molecular Parameters

Figure5 shows the molecular orbital representation of the H2L ligand and its Cd(II) complex along with their HOMO, LUMO energies and energy band gaps. Both the highest occupied molecular orbital (HOMO) and lowest unoccupied molecular orbital (LUMO) were the main orbitals that participate in chemical stability [11, 44, 45]. The HOMO represented the ability to donate an electron; LUMO as an electron acceptor represents the ability to obtain an electron. In other words, the energy of the HOMO is directly related to the ionization potential; LUMO energy is directly related to the electron affinity. Hence, these frontier energies play an important role in the electric and optical properties. The energy gap between the HOMO and LUMO energies is a critical parameter in determining molecular electrical transport properties due to providing measurement of electron conductivity. In addition, this energy gap characterizes the molecular stability and spectroscopic properties of the molecular systems. The smaller energy gap describes a chemically soft molecule which can be easily polarizable. The HOMO-LUMO energy gap has proved to be an important tool in determining the kinetic stability and chemical reactivity of a molecule. The lower energy gap is a suitable condition where a molecule can be exited easily, whereas a higher energy gap leads to higher kinetic stability but lower chemical reactivity of the molecule. The difference between the energy gap for the ligand reflects the presence of the complexation status. The increase in the global electrophilicity value attributed to the highest capacity of accepted electrons, so the ligand has a highly powerful donation ability. The calculation of both absolute hardness (η) and absolute softness (σ) parameters is useful to recognize the molecular stability and reactivity. The calculations of the binding energy indicated that the increase of the value of the calculated binding energy of the complex compared to that of the free ligand indicated that the stability of the formed Cd(II) complex was higher than that of the free diazo ligand.Figure 5
Surface phase of frontier orbitals of H2L and [Cd(H2L)Cl2].H2O.Additional parameters such as chemical potentialsPi, global electrophilicity ω, global softness S, electrophilicity index (χ), and additional electronic charge ∆Nmax were calculated for the free H2L ligand and the Cd-H2L complex (Table 3). The high ω value of both free H2L and Cd(II) complex suggested a great chance and priority for biological activity which is further confirmed by the experimental data.Table 3
The calculated quantum chemical parameters of H2L and its Cd(II) complex.
The quantum parameterH2L[Cd(H2L)Cl2].H2OE (a.u.)−1402.46−140.88Dipole moment (Debye)1.3718.022EHOMO (eV)−6.095−5.61ELUMO (eV)−2.977−4.096ΔE (eV)3.3491.508`χ(eV)4.5344.853η (eV)1.5590.757σ (eV)−10.6411.321Pi (eV)−4.534−4.853S (eV)−10.3210.661ω (eV)6.59315.559∆Nmax2.96.412The mentioned quantum chemical parameters were calculated with the help of the following equations [43, 44]:(4)ΔE=ELUMO−EHOMO,χ=−EHOMO+ELUMO2,η=ELUMO−EHOMO2,σ=1η,Pi=−χ,S=12η,ω=Pi22η,ΔNmax=Piη.


## 3.7. Vibrational Properties

In order to acquire the spectroscopic signature of the diazo ligand and its complexes, a frequency calculation analysis was performed. Vibrational properties had been investigated to determine vibrational modes connected with the molecular structure of the cadmium (II) complex using FT-IR spectra based on the optimized geometry. It is well known that the vibrational frequencies computed at quantum chemical methods such as DFT levels contain well-known systematic errors. The scaling factor of 0.96 for the LanL2DZ level was used to correct the effects of anharmonicity and neglected part of electron correlation [11, 44, 45]. The theoretical and the experimental spectra are shown in Supplementary Figure 8. It is also noteworthy that experimental results belong to the solid phase, while theoretical calculations belong to a gas phase.The stretching vibration of a free or nonhydrogen-bonded OH group of H2L appears at 3664 cm−1 [11, 45]. υ(COO−)asy and υ(COO−)sy stretching of the free ligand appeared at 1597 cm−1 and 1392 cm−1, respectively. The theoretical spectrum of the free azo ligand showed υNH2 and υN = N modes at 3637 and 1694 cm−1, respectively. In the Cd(II) complex, the chelating ring occurred by means of the ligand bonding to the cadmium (II) ion with a tetradentate N,O chelated mode. It was observed from Figures 6 that the wavenumber interval between the υ(COO−)asy and υ(COO −)sy stretching modes is an effective indicator of forming complexation. The wavenumber value of υN = N for the complex was lower than free H2L which was an indicator of the coordination through the nitrogen atom. The spectrum of the Cd(II) complex showed υNH2 at a wave number very close to free H2L at 3639 cm−1 confirming that this group did not participate in the complex formation process. υM-O, υM-N, and υM-Cl modes of the Cd(II) complex were found at 564, 487, and 326 cm−1, respectively. The calculated and experimental data were very close to each other (Supplementary Figure 8).Figure 6
The proposed structures of H2L and its complexes.


## 3.8. Structural Interpretation of Metal Complexes

The structures of the complexes of the H2L diazo ligand with Cr(III), Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) ions were conﬁrmed by the elemental analyses, IR, molar conductance, magnetic, mass, solid reﬂectance, and thermal analysis data. Therefore, from the IR spectra, it was concluded that H2L behaved as a neutral tetradentate ligand coordinated to the metal ions via two nitrogen atoms of azo groups (N=N) and two oxygen atoms of protonated carboxylate groups (COO). From the molar conductance data, it was found that the complexes were electrolytes. On the basis of the above observations and from the magnetic and solid reﬂectance measurements, octahedral geometry was suggested for the investigated complexes (Figure 6).


## 3.9. Antimicrobial Activity of Metal Complexes

### 3.9.1. Antifungal Activities

The preliminary fungi toxicity screening of the H2L diazo ligand and complexes was performed against Candida albicans in vitro by the diffusion technique [46]. H2L free ligand and Cr(III), Mn(II), and Fe(III) complexes showed no fungal growth inhibition, but Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) complexes showed fungal growth inhibition at different rates (Figure 7). The Cd(II) and Ni(II) complexes in this study were nearly two times more active than the ketokonazole standard against the Candida albicans microorganism used (Table 4).Figure 7
Biological activity of H2L and its complexes.Table 4
Biological activity of the H2L ligand and its metal complexes.
SampleInhibition zone diameter (mm/mg sample)Bacillus subtilisEscherichia coliNeisseria gonorrhoeaeStaphylococcus aureusCandida albicansControl: DMSO0.00.00.00.00.0Ligand (H2L)0.00.00.00.00.0[Cr(H2L)Cl2]Cl.2H2O141515150.0[Mn(H2L)Cl2]91090.00.0[Fe(H2L)Cl2]Cl.H2O9990.00.0[Co(H2L) Cl2]1616151612[Ni(H2L)Cl2]1414151316[Cu(H2L)Cl2].3/2H2O1212121310[Zn(H2L)Cl2].H2O111213159[Cd(H2L)Cl2].H2O1815132317Amikacin69760Ketokonazole————9


### 3.9.2. Antibacterial Activities

The antibacterial activities of the H2L diazo ligand and its complexes against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli are presented in Table 4. The H2L ligand was found to have no activity at all towards both different types of bacteria (Gram-positive Bacillus subtilis and Staphylococcus aureus and Gram-negative Neisseria gonorrhoeae and Escherichia coli). This can be attributed to their very versatile nutritional capability, adaptability to various hydrocarbon rings, and the possession of the pump mechanism which ejects metal complexes as soon as they enter the cell [47]. In addition, Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli were sensitive to all the complexes, and an inhibitory zone was obtained in the range of 9.0–23.0 mm (Figure 7).In all cases, the metal complexes were found to be more active than the H2L ligand expectedly due to chelation, which reduced the polarity of the metal atom, mainly because of partial sharing of its positive charge with donor groups of the ligand and possible π-electron delocalization on the aromatic rings. This increased the lipophilic character, favouring its permeation into the bacterial membrane, causing the death of the organisms [48].Higher activity of the metal complex was probably due to greater lipophilic nature of the complex. It increased the activity of the metal complex and can be explained on the basis of Overtone’s concept and Tweedy’s chelation theory [49–52]. According to Overtone’s concept of cell permeability, the lipid membrane that surrounds the cell favours the passage of only lipid-soluble materials due to which liposolubility was considered to be an important factor that controls the antimicrobial activity. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups [53, 54]. Furthermore, it increased the delocalization of the π electrons over the whole chelate ring and enhanced the lipophilicity of the complex. This increased lipophilicity enhanced the penetration of the complexes into the lipid membrane and thus blocked the metal-binding sites on enzymes of microorganisms [55]. These metal complexes also disturb the respiration process of the cell and thus block the synthesis of proteins, which restricts further growth of the organism [56].A look at the antibiotic, amikacin, activities (6.0–9.0 mm) against various bacterial isolates relative to the metal complexes (9.0–23.0 mm) showed that the activities of the former were much lower, with optimum activity being about half of metal complexes, against all the bacterial organisms.When the antimicrobial activity of metal complexes is investigated, the following principal factors [57] should be considered: (i) the chelate eﬀect of the ligands; (ii) the nature of the N-donor ligands; (iii) the total charge of the complex; (iv) the existence and the nature of the ion neutralizing the ionic complex; and (v) the nuclearity of the metal center in the complex. This is probably one of the reasons for the diverse antibacterial activity shown by the complexes, while the nature of the metal ion coordinated to the H2L ligand may have a signiﬁcant role to this diversity. In general, all the complexes exhibited better inhibition than free H2L against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli (Table 4). More speciﬁcally, Cd(II) and Co(II) complexes showed the best inhibition among all the complexes in this study, and they were one to twenty-three times more active than the H2L diazo ligand against all the microorganisms used, indicating that the coordination of the H2L ligand to Cd(II) and Co(II) ions enhanced its antimicrobial activity (Figure 7). On the contrary, the rest complexes present higher antimicrobial activity to H2L against the four microorganisms.


## 3.9.1. Antifungal Activities

The preliminary fungi toxicity screening of the H2L diazo ligand and complexes was performed against Candida albicans in vitro by the diffusion technique [46]. H2L free ligand and Cr(III), Mn(II), and Fe(III) complexes showed no fungal growth inhibition, but Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) complexes showed fungal growth inhibition at different rates (Figure 7). The Cd(II) and Ni(II) complexes in this study were nearly two times more active than the ketokonazole standard against the Candida albicans microorganism used (Table 4).Figure 7
Biological activity of H2L and its complexes.Table 4
Biological activity of the H2L ligand and its metal complexes.
SampleInhibition zone diameter (mm/mg sample)Bacillus subtilisEscherichia coliNeisseria gonorrhoeaeStaphylococcus aureusCandida albicansControl: DMSO0.00.00.00.00.0Ligand (H2L)0.00.00.00.00.0[Cr(H2L)Cl2]Cl.2H2O141515150.0[Mn(H2L)Cl2]91090.00.0[Fe(H2L)Cl2]Cl.H2O9990.00.0[Co(H2L) Cl2]1616151612[Ni(H2L)Cl2]1414151316[Cu(H2L)Cl2].3/2H2O1212121310[Zn(H2L)Cl2].H2O111213159[Cd(H2L)Cl2].H2O1815132317Amikacin69760Ketokonazole————9


## 3.9.2. Antibacterial Activities

The antibacterial activities of the H2L diazo ligand and its complexes against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli are presented in Table 4. The H2L ligand was found to have no activity at all towards both different types of bacteria (Gram-positive Bacillus subtilis and Staphylococcus aureus and Gram-negative Neisseria gonorrhoeae and Escherichia coli). This can be attributed to their very versatile nutritional capability, adaptability to various hydrocarbon rings, and the possession of the pump mechanism which ejects metal complexes as soon as they enter the cell [47]. In addition, Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli were sensitive to all the complexes, and an inhibitory zone was obtained in the range of 9.0–23.0 mm (Figure 7).In all cases, the metal complexes were found to be more active than the H2L ligand expectedly due to chelation, which reduced the polarity of the metal atom, mainly because of partial sharing of its positive charge with donor groups of the ligand and possible π-electron delocalization on the aromatic rings. This increased the lipophilic character, favouring its permeation into the bacterial membrane, causing the death of the organisms [48].Higher activity of the metal complex was probably due to greater lipophilic nature of the complex. It increased the activity of the metal complex and can be explained on the basis of Overtone’s concept and Tweedy’s chelation theory [49–52]. According to Overtone’s concept of cell permeability, the lipid membrane that surrounds the cell favours the passage of only lipid-soluble materials due to which liposolubility was considered to be an important factor that controls the antimicrobial activity. On chelation, the polarity of the metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor groups [53, 54]. Furthermore, it increased the delocalization of the π electrons over the whole chelate ring and enhanced the lipophilicity of the complex. This increased lipophilicity enhanced the penetration of the complexes into the lipid membrane and thus blocked the metal-binding sites on enzymes of microorganisms [55]. These metal complexes also disturb the respiration process of the cell and thus block the synthesis of proteins, which restricts further growth of the organism [56].A look at the antibiotic, amikacin, activities (6.0–9.0 mm) against various bacterial isolates relative to the metal complexes (9.0–23.0 mm) showed that the activities of the former were much lower, with optimum activity being about half of metal complexes, against all the bacterial organisms.When the antimicrobial activity of metal complexes is investigated, the following principal factors [57] should be considered: (i) the chelate eﬀect of the ligands; (ii) the nature of the N-donor ligands; (iii) the total charge of the complex; (iv) the existence and the nature of the ion neutralizing the ionic complex; and (v) the nuclearity of the metal center in the complex. This is probably one of the reasons for the diverse antibacterial activity shown by the complexes, while the nature of the metal ion coordinated to the H2L ligand may have a signiﬁcant role to this diversity. In general, all the complexes exhibited better inhibition than free H2L against Bacillus subtilis, Staphylococcus aureus, Neisseria gonorrhoeae, and Escherichia coli (Table 4). More speciﬁcally, Cd(II) and Co(II) complexes showed the best inhibition among all the complexes in this study, and they were one to twenty-three times more active than the H2L diazo ligand against all the microorganisms used, indicating that the coordination of the H2L ligand to Cd(II) and Co(II) ions enhanced its antimicrobial activity (Figure 7). On the contrary, the rest complexes present higher antimicrobial activity to H2L against the four microorganisms.


## 3.10. Molecular Docking

Targeting the minor groove of DNA through binding to a small molecule has long been considered an important tool in molecular recognition of a speciﬁc DNA sequence. DNA is considered a major biological target for metal complexes that have pharmacological activity. In this context, molecular docking between H2L and [Cd(H2L)Cl2].H2O with four possible biological targets, receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and the crystal structure of Escherichia coli (PDB ID: 3T88), was performed. The structures of these macromolecules were obtained from the Protein Data Bank. The calculated binding energies for the interaction with these proteins are listed in Tables 5 and 6 . The 3D plot curves of docking with free H2L and [Cd(H2L)Cl2].H2O are shown in Figures 8 and 9.Table 5
Energy values obtained in docking calculations of H2L with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88).
ReceptorLigand moietyReceptorInteractionDistanceE (kcal/mol)5JPEN 41O HIS 413 (A)H-donor3.07−2.5N 14NE2 HIS 290 (A)H-acceptor3.33−2.3O 34NH2 ARG 386 (A)H-acceptor3.01−2.8O 36NH2 ARG 261 (A)H-acceptor2.93−1.4O 36NE2 HIS 290 (A)H-acceptor2.99−1.13HB5N 41OD1 ASN 152(X)H-donor2.83−3.1N 42O GLY 9(X)H-donor3.44−0.7O 36OG SER 142(X)H-acceptor2.87−1.56-ringCD2 PHE 192(X)Pi-H3.75−0.63Q8UO 38OE1 GLU 95 (A)H-donor2.90−6.0N 42OE2 GLU 51 (A)H-donor3.00−2.0N 14N THR 91 (A)H-acceptor3.54−1.6O 34NZ LYS 9 (A)H-acceptor3.58−0.6O 34NH2 ARG 102 (A)H-acceptor3.04−2.9O 38OE1 GLU 95 (A)H-donor2.90−6.03T88O 38O PHE 162 (A)H-donor3.33−0.6N 41OD2 ASP 163 (A)H-donor3.11−2.2N 14OH TYR 97 (A)H-acceptor3.18−1.2O 36OH TYR 97 (A)H-acceptor3.38−0.6Table 6
Energy values obtained in docking calculations of [Cd(H2L)Cl2].H2O with receptors of the crystal structure of S. aureus (PDB ID: 3Q8U), crystal structure of protein phosphatase (PPZ1) of Candida albicans (PDB ID: 5JPE), receptors of breast cancer mutant oxidoreductase (PDB ID: 3HB5), and crystal structure of Escherichia coli (PDB ID: 3T88).
ReceptorLigand moietyReceptorInteractionDistanceE (kcal/mol)5JPEN 42OD2 ASP 373 (A)H-donor2.89−3.1O 48O ARG 386 (A)H-donor3.07−0.7O 38NH2 ARG 386 (A)H-acceptor3.16−2.46-ringCE2 TYR 437 (A)Pi-H4.62−0.63HB5O 48O CYS 185 (X)H-donor2.53−0.2O 48OG SER 142 (X)H-acceptor2.68−4.76-ringCG1 VAL 225 (X)Pi-H4.14−0.73Q8UN 10OE1 GLU 51 (A)H-donor2.72−18N 10OE2 GLU 51 (A)H-donor3.38−0.7CL 36O HIS 48 (A)H-donor3.68−0.7O 48O GLY 116 (A)H-donor2.79−2.2O 50NH1 ARG 125 (A)H-acceptor2.85−4.4N 10OE1 GLU 51 (A)Ionic2.72−6.7N 10OE2 GLU 51 (A)Ionic3.38−2.4N 10OD1 ASP 118 (A)Ionic2.69−7.03T88N 10O PHE 162 (A)H-donor2.89−5.2O 38OH TYR 97 (A)H-acceptor2.68−2.0Figure 8
3D plot of the interaction between the H2L ligand with receptors of (a) 5JPE, (b) 3HB5, (c) 3Q8U, and (d) 3T88.
(a)(b)(c)(d)Figure 9
3D plot of the interaction between Cd-H2L with receptors of (a) 5JPE, (b) 3HB5, (c) 3Q8U, and (d) 3T88.
(a)(b)(c)(d)From these data, the following can be concluded [58–60]:(1)
First, H2L and its Cd(II) complex showed strong binding affinities for all proteins(2)
The [Cd(H2L)Cl2].H2O complex binds stronger than the free ligand to different proteins, see Figure 10(3)
The binding energy for the interaction of the cadmium (II) complex with the receptor protein 3Q8U was three times less than the free diazo ligand confirming the high activity of this compound [58](4)
A closer look at the interactions indicated that the presence of heteroatoms such as oxygen and nitrogen was critical to how the compound binds [59](5)
Generally speaking, H-acceptor interaction dominates the binding of the synthesized complexes to different receptors besides the metal binding to the protein backboneFigure 10
The binding energy of the interaction between H2L and Cd-H2L and receptors of 5JPE, 3HB5, 3Q8U, and 3T88.


## 4. Conclusion

A new diazo ligand, 2,2'-(1,3-phenylenebis(diazene-2,1-diyl))bis(4-aminobenzoic acid), was synthesized and treated with different metal salts to give the corresponding metal complexes. The new ligand H2L reacted with the metal ions as 1 : 1 unit. The analytical data showed that reactions of the free diazo ligand (H2L) with Mn(II), Co(II), Ni(II), Cu(II), Zn(II), and Cd(II) formed complexes with chemical formulae [M(H2L)Cl2].nH2O (M = Co(II), Ni(II), Cu(II), Mn(II), Zn(II) and Cd(II)) and [M(H2L)Cl2]Cl.nH2O (M = Cr(III) and Fe(III)). The possible structures of the ligand and its metal-azo complexes were proposed based on elemental analyses, 1H NMR, ESR, MS, IR, and UV-Vis electronic absorption. It was concluded that the H2L ligand behaved as a neutral tetradentate ligand with N coordination sites and protonated carboxylate oxygen upon coordination to the metal ions leading to the formation of octahedral geometries for all complexes. The thermal studies for the compounds showed a higher thermal stability for the complexes than that of the free ligand. This may be due to chelation, the type, and the number of solvents of crystallization in the metal complexes. Also, the ligand and its metal complexes were screened in vitro against microorganisms (bacteria or fungi). The results of biological activity showed that the Cd(II) complex had higher antibacterial activity against the (Gram-positive) bacteria and fungi, while Co(II) had higher antibacterial activity against the (Gram-negative) bacteria compared to the free ligand and other metal complexes. Molecular docking studies of the free ligand and the Cd(II) complex with four different receptors clarified that the Cd(II) complex showed the minimum binding ability with the binding energy of −18 kcal/mol with the crystal structure of S. aureus (3Q8U).



---

*Source: 6382037-2020-10-23.xml*