Solvent free synthesis, characterization, DFT, cyclic voltammetry and biological assay of Cu(II), Hg(II) and UO₂(II) – Schiff base complexes

2.1 Chemicals

The compounds utilized to prepare Schiff base ligand and its metal ion complexes were 3-hydrazineyl-3-oxo-N-phenylpropanamide, isatin, Cu(OAc)₂·2H₂O(98%), UO₂(OAc)₂·2H₂O (98%) and HgCl₂ (99%). Absolute ethanol (EtOH) and dimethyl sulfoxide (DMSO) solvents employed either for the ligand synthesis or to dissolve the complexes for analytical purposes, were spectroscopic in nature. The reagents of 2, 2′-azinobis-(3-ethylbenzthiazoline-6-sulphonic acid)(ABTS) and L-ascorbic acid which utilized for biological tests were provided by Sigma-Aldrich. Also, HepG2 cell line was donated by ATCC's company which holding biological products and vaccines (VACSERA) for cytotoxic assay. Notice: all these reagents were purchased from Sigma-Aldrich were BDH in nature.

2.2 Synthesis of the ligand and its metal complexes

In 50 mL absolute ethanol, equal moles (1:1) of 3-hydrazineyl-3-oxo-N-phenylpropanamide (0.01 mol; 1.93 g) and isatin (0.01 mol; 1.47 g) were mixed to isolate the ligand. For 3 h, the reaction mixture was maintained at 80 °C during reflux. Then after that the mixture was concentrated to reach to the half volume, and set aside to cool. Filtration was used to extract the product, which was then recrystallized from EtOH and dried in vacuum desiccators over anhydrous CaCl₂. The ligand isolated was (E)-3-oxo-N'-(2-oxoindolin-3ylidene)-3-(phenylamino) propane hydrazide and abbreviated H₃L according to existence of three labile hydrogen atoms that may be ionization easily (Scheme 1). Equi-molar ratios from the ligand (3 mmol, 0.966 g) and each of Cu(OAc)₂·2H₂O (3 mmol, 0.653 g), UO₂(OAc)₂·2H₂O (3 mmol, 1.272 g) or HgCl₂ (3 mmol, 0.815 g) salt were grinded without unpleasant solvent except few drops of water to aid grinding, only. The time taken to prepare the Cu(II) complex is smaller (5 min) than that used with the other complexes (15 min). The acidic smell was noticed during the synthesis of Hg(II) and UO₂(II) complexes, which denotes the deprotonation of the ligand due to high heat generated during vigorous grinding inside ball milling. The first indicator of the reaction is the color of the solid resulted, which is entirely different from the reactants and thin layer chromatography(TLC) test confirms the product purity (Antony et al., 2013; Al-Qahtani et al 2012; Munshi et al., 2021;). Notice: Table 1, lists some physical and analytical properties of the new compounds.

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Scheme 1

Synthesis outlines for the ligand and its metal complexes.

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2.3 Instrumentation

For clarity, the approaches employed for the analyses are depicted as images in the supporting materials file (Scheme S1). Moreover, each device's description has been given below, while the implementation conditions will be reported individually in the discussion section. Complexometric titration was used to assess the metal content (without uranium), while the gravimetric analysis was used to determine the chloride content (Vogel et al., 1989). On the other side, after complete combustion of weighed sample in open air and then in oven at 800 °C until the organic compound was completely evaporated, the percentage of uranium was calculated from the residual part weight (U₃O₈).

2.4 Geometry optimization

Gaussian 09 (Frisch et al., 2009) was deemed the best molecular modeling application that was used to verify practically accepted structures for the substances under investigation. The DFT/B3LYP method that implemented the valent electron basis set (SDD) (Lee et al., 1988; Li et al., 2016) and the polarized function within valence double-zeta basis set (6-31G*), was applied for optimization [Li et al., 2016]. The SDD was identified as the best core potential basis set for the Hg(II) complex [Li et al., 2016], while 6-31G* was sufficient for organic ligands and the other complexes, with 6-d-functions added as polarized functions to 6-31G set. The DFT was done at B3LYP level (Becke et al 1993) to investigate both states of ground and excited attributes utilizing Integral Equation Formalism Variant (IEF-PCM) with a polarizable continuum model. Using Gauss-View and Gauss-Sum 2.2 programs, all quantum data was extracted from two exported computational files, as well as the third that extracted after formulation for chk file on Gauss-prog (fchk) (Dennington et al., 2007). Finally, we were unable to apply an appropriate basis set for structural optimization of UO₂(II) complex since we couldn't discover its computational properties.

2.5 Cyclic voltammetry

This electro-analytical technique was applied for CuCl₂ in solution either in absence or in presence of the ligand. The cell used is a small beaker containing 0.L M KCl (30 mL) as a supporting electrolyte. A reference electrode, Ag/AgCl/KCl (saturated), the glassy carbon working electrode (GCWE), and platinum wire as an auxiliary electrode were used. These electrodes were dipped in KCl solution and connected with potentiostat.

2.6 Biological activity

3.1 Analytical features

To explain their chemical formulae, the ligand (H₃L) and its corresponding Cu(II), Hg(II), and UO₂(II) complexes, were physically and chemically studied (Table 1). The new complexes appeared with high stability, non-moisture absorbent, and high melting points. The non-conducting property of all complexes was assumed based on the conductivity of 10^-3M (Λ_m, Ω^-1cm²mol⁻¹) that determined in DMSO solvent (Geary et al 1971). All complexes were recommended in equal molar ratios (1 M:1L) based on elemental analysis, which was later confirmed by EDX analysis (part 3.5).

3.2 Spectral analysis for Schiff base ligand

The most notable infrared bands of H₃L and its complexes are mentioned in Table 2. The bands found in the IR spectrum of H₃L (Fig. S1) at 1742, 1686, and 1662 cm⁻¹ are attributing to υ(C = O)₁, υ(C = O)₂ and υ(C = O)₃ (Alharbi et al., 2021) vibrations, respectively. A weak band at 1078 cm⁻¹ is attributing to υ(N-N) vibration. Sequentially, the vibrations of υ(NH)₁, υ(NH)₂, υ(NH)₃ and υ(CH)_aromatic, were appeared at 3200, 3303, 3259, and 3009 cm⁻¹ (Alharbi et al., 2021; Nakamoto, 1970). The band at 3451 cm⁻¹ may assign to υ(C-OH), which produced via tautomerism of amide groups (NH-C = O). This feature reveals the mixed presence for the tautomer forms (keto/enol). Surprisingly, molecular modeling using DFT/B3LYP approach verifies the comparable stability of the two forms (keto/enol), implying that their mixed presence produces equivalent probability of their coordination (see part 2.7).

The ¹H NMR (DMSO‑d₆) spectrum of the ligand displayed signals supporting its tautomeric forms as follow; δ 3.82 (s, 2H, CH₂), 6.94–7.81 (m, 8H, Ar-H), 8.12 (d, J = 7.40 Hz, 1H, Ar-H), 10.30 (s, 1H, N(2)H), 10.78 (s, 1H, N(3)H), 11.30 (s, 1H, N(1)H), 12.52 ppm (s, 1H, OH) (Fig. S2). Moreover, the mass spectrum (Fig. S3) of the ligand displayed the molecular ion peak for [M]^.+at m/z = 323.24 (0.87 %), which is comparable to its molecular weight and corresponds to C₁₇H₁₄N₄O₃ formula. The fragmentation path of the ligand under influence of bombarded electrons was suggested in Scheme 2.

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Scheme 2

The pattern of fragmentation for H₃L ligand 3.3. IR, ¹HNMR and mass spectra of the complexes.

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In [Cu(H₃L)(OAc)₂].H₂O complex spectrum, H₃L functions as a neutral tetradentate ligand which coordinates via (C = O)₁,(C = O)_2, (C = O)₃ and azomethine nitrogen (C = N)_azo. The movement of υ(C = O)₁, υ(C = O)₂, υ(C = O)₃ and υ(C = N) to lower wavenumbers aided this coordination mode (Table 2). The fact that υ(NH)_1, υ(NH)₂ and υ(NH)₃ were shifted to higher wavenumbers, suggests that they are no longer coordinated. Furthermore, the novel bands detected at 511 and 459 cm^-1can be attributed to υ(Cu-O) and υ(Cu-N), respectively (Nakamoto et al., 1970). The Cu(II) complex has two bands at 1547 and 1344 cm⁻¹, correspond to acetate group vibrations; υ_s(O-C-O) and υ_as(O-C-O). The monodentate character of the acetate group was indicated from the difference between those two bands (203 cm⁻¹).

H₃L acts as a mononegative tetradentate within [UO₂(H₂L)(OAc)] and [Hg(H₂L)(H₂O)Cl] complexes. The two carbonyl oxygen [(C = O)₁& (C = O)₂], the deprotonated enolized carbonyl group (C-O^-)₃ and the azomethine nitrogen (C = N)_azo were the coordinating sites within the two complexes. This chelation mode is based on the simultaneous development of new bands at 1632; 1622 and 1011; 1131 cm⁻¹, which are assigned to υ(C = N*)₃ and υ(C-O)_enolic(3) (Alharbi et al., 2021), respectively, as well as the disappearance of both υ(C = O)₃ and υ(NH)₃ vibrations. Also, this mode was supported by the movement of υ(C = O)₁, υ(C = O)₂ and υ(C = N) vibrations to lower wavenumbers. The vibrations of υ(M−O) and υ(M−N) (Alharbi et al., 2021) were detected at 571; 476 and 530; 467 cm⁻¹, respectively. The UO₂(II) complex spectrum displayed two bands at 1547 and 1457 cm⁻¹, which correspond to υ_s(O-C-O) and υ_as(O-C-O) vibrations in acetate groups. The difference between those bands (90 cm⁻¹) denotes the acetate group's bidentate nature (Nakamoto et al., 1970). The lack of (NH)₃ and (OH) signals in the ¹H NMR spectra of the two diamagnetic complexes (UO₂(II) & Hg(II)) underlines the deprotonation of enolized carbonyl oxygen (=C-O-)₃ (i.e., Fig. S2). The ¹H NMR spectrum of Hg(II) complex displayed the following signals; δ 3.93 (s, 2H, CH₂), 7.13 (d, J = 8.40 Hz, 2H), 7.34 (t, J = 7.20 Hz, 1H), 7.48–7.752 (m, 4H), 7.97 (d, J = 7.80 Hz, 1H), 8.08 (d, J = 7.80 Hz, 1H), 11.82 (s, 1H, NH) and 11.37 ppm (s, 1H, NH). Two sharp bands at 856–868 and 540–581 cm^-1regions are attributed to ρ_r(H₂O) and ρ_w(H₂O) (Nakamoto et al., 1970) vibrations, respectively, for coordinating water molecule in Hg(II) complex. Furthermore, the [UO₂(H₂L)(OAc)] complex's IR spectrum shows three bands at 920, 870, and 268 cm⁻¹, which correspond to υ₃, υ₁ and υ₄ vibrations of dioxouranium ion, respectively. Based on McGlynn et al technique (McGlynn et al., 1961), the υ₃ value was utilized to determine the force constant (F) of υ(U = O) from (υ₃)²= (1307)²(F_U-O)/14.103. The uranyl complex force constant was then used to calculate the U-O bond length (by Å) using Jones's equation: R_U-O = 1.08 (F_U-O)^-⅓ +1.17 (McGlynn et al., 1961). The computed F_U-O and R_U-O values are 6.987 mdynes Å⁻¹ and 1.735 Å, respectively, which are within the uranyl complexes' normal range. Thus, the neutral tetradentate mode of bonding was recorded with Cu(II) ion while the mononegative tetradentate mode was recorded with UO₂(II) or Hg(II) ion. This may refer to the sizes of Cu(II), Hg(II) and U(VI) ions as 1.57, 1.02 and 0.66 Å, respectively. So, the metals with smaller sizes (Hg(II) and U(VI)) prefer the covalence attachment which happened after the ionization of amide group. Also, the time used for grinding the reactants to synthesize the complexes, effects on the mode of bonding. The long time consumed in grinding to prepare UO₂(II) and Hg(II) complexes leads to high heat that may activate the ionization of the ligand particularly with presence of anions (acetate, chloride) in the reaction medium.

3.3 Magnetism, UV–Vis and ESR spectra

The [Cu(H₃L)(OAc)₂].H₂O complex revealed a magnetic moment value of 2.1B.M., which is consistent with that known for d⁹-systems. Also, in tetragonally deformed octahedral shape, its spectrum (Fig. S4) reveals a band at 21,786 cm⁻¹ and a broad band at 15479 cm⁻¹, which may assign to ²B_1g→²E_g and ²E_g→²A_1g transitions, respectively (Lever et al 1984). The UV–Vis spectrum of [UO₂(H₂L)(OAc)] complex reveals two bands at 29,762 and 23148 cm⁻¹, which attributing to ¹Σ⁺_g→ ²π _u transition in dioxouranium (VI) and n → π* charge transfer, respectively (Al-Fahemi et al 2017).

At room temperature, ESR spectrum of Cu(II) complex at υ = 9.435 GHz (Fig. 1) exhibited axially symmetric g factor with g_‖ >g_⊥>2.0023 which agrees with ²A₁g (d_X²-_Y²), as a ground term for octahedral geometry (Abou-Hussen et al., 2005; Hathaway et al., 1984). The G factor could being calculated by using the following expression: G = (g_‖– 2) / (g_⊥–2) = 4. The factor G as the exchange interaction parameter (Welleman et al., 1978), if it is bigger than 4, the exchange interaction between copper centers in the solid sample is insignificant, according to Hathaway. But, if it<4, the exchange interaction between copper centers in the solid sample is very significant. The computed value (4.6) indicates that there is no interaction between copper centers in neighboring molecules. Also, the g_‖/A_‖ ratio was 138, which indicates that the deformed octahedral has a critical distortion for dihedral angle in xy-plane. The values 2.28, 2.06, and 165 cm⁻¹ were determined as the g-tensor parameters of g_‖, g_⊥ and A_‖, respectively. The degree of covalence of either in plane σ-bonding or in-plane π-bonding between copper ion and the ligand, was estimated by using known equations (Bagdatil et al., 2017). The values of α² and β² appear to be 0.80 and 0.77, respectively. These values revealed the dominance of ionic property in Cu(II)-L bonds.

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Fig. 1

ESR spectrum of [Cu(H₃L)(OAc)₂].H₂O Complex.

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3.4 SEM, XRD and EDX studies

The surface topography of various complexes (Fig. S5) was studied using accelerated beam of electrons in SEM technique (Al-Fahemi et al., 2018). As seen, a regular spherical particles are randomly dispersed between particles produced as frozen rocks in the surface of UO₂(II) and Hg(II) complexes. On the other side, the Cu(II) complex morphology incorporates spherical particles in the form of shattered glacial boulders and micrometer-sized particles.

Cu/Ka derived X-ray diffraction (λ = 1.5406) which directed to investigate the crystal lattice of the Cu(II) complex at 25 °C. The diffraction pattern of the investigated compound (Fig. 2) was extracted in the range of 10° < 2θ < 80°(Khan et al., 2013). Sharp peaks noticed in the model, reveal a well-defined crystal structure formed. The crystal size was calculated using Debye-Scherrier equation [b = 0.94 λ/(S cos θ)] based on FWHM (b) for the characteristic peak. The crystal lattice parameters and particle size had magnitudes of 21.94, 2.062, 0.23, and 530, respectively, corresponding to θ^o, d (A^o), FWHM (β) and I. The particle size S (nm) was found to be 1.47 nm and based on Bragg equation (nλ = 2d sin θ) at n = 1, the inner crystal plane d-spacing value was calculated (1.945 Å).

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Fig. 2

XRD pattern of [Cu(H₃L)(OAc)₂].H₂O complex.

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Using energy-dispersive x-ray microanalysis (Alkhatib et al 20 20), the complex formation was confirmed, and the elemental mass percentage (wt %) was calculated. An induced radiation (x-ray) was directed towards the solid samples to access inner shells (K- or L-shell) and interact with resident electrons. Following this, the electrons are evacuated and then creating positively charged gaps which can be filled with electrons from L or M shells. K-series (Ka1, Ka2, or Kb) or L-series was created from energy loss according to the shift from upper levels (L or M) to lower level (Ka1, Ka2, or Kb). The L-series, which is useful in EDX analysis, can be used to get the elemental percentage. Each element's unique electrical structure is represented by a singlet-set of peaks (Fig. 3). The heavier element (metal) was employed to express mass percentage (wt%), while lighter elements (C, H, and O) were not. This identification was carried out over a long time to improve the signals in a variety of locations. The elemental molar ratios computed (wt/At%) matched those obtained from analyses of elements (C, N, O, Cl & M).

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Fig. 3

EDX patterns of(A) UO₂(II), (B) Hg(II) and (C) Cu(II) complexes.

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3.5 TGA of Cu(II) complex

In the temperature range of 37–125 °C, the TGA curve of [Cu(H₃L)(OAc)₂].H₂O complex (Fig. S6) exhibits 3.44 % weight loss, which corresponds to water of crystallization. The TGA curve shows 23.08 % weight loss between125 and 258 °C, which could assign to the elimination of two acetic acid molecules (Rakha et al., 1989). The TGA curve shows a 17.50% weight loss in the range of 285–430 °C, which may be due to a weakly bonded moiety (C₆H₅N). The residual part (CuC₁₁H₆O₃N₃) was recorded at 430 °C with weight percentage of 55.98. Such proposed behavior over the TGA curve, confirms the existence of crystal water molecule and the whole molecular formula of the complex (Table S1). While, the incomplete degradation for the complex indicates the stability of coordination environment surrounds the copper ion.

3.6 Mass spectral analysis

In the lack of chance to isolate x-ray single crystal, the mass spectral analysis has necessity to confirm the molecular formulae of the new complexes which suggested based on former studies (Fig. S7). The signals yielded from fragmentation of the two selected complexes [UO₂(II) and Hg(II)] were obtained over a broad mass range (m/z = 50–1000) under a fixed rate of heating (40 °C/min) and at 70 ev. The molecular ion peaks of UO₂(II) and Hg(II) complexes were appeared at m/z = 651.83(calcd. 649.34) and m/z = 558.85 (calcd. 574.44). The molecular ion peaks agree perfectly with the molecular weights suggested from elemental analysis but with Hg(II) complex the recorded peak assign to [(M⁺+1) –H₂O]

3.7 Structural verification via computations

The DFT/B3LYP approach was used to improve the structures of the ligand forms (keto/enol) and two of its complexes [Cu(II) & Hg(II)] using Gaussian 09 software (Frisch et al., 2009). All compounds were optimized under 6-31G* basis set except the Hg(II) complex, which was optimized by using the efficient core potential basis set (SDD) level (Lee et al., 1988; Li et al., 2016). The best structural forms (Fig. S8), frontier orbital patterns (Fig. 4), and electrostatic potential (MEP) were extracted from produced files (Fig. S9).

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Fig. 4

the patterns of HOMO and LUMO levels in the optimized structures.

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The optimized geometry of H₃L ligand revealed a specific arrangement for C(8) = O(13), C(10) = O(14), N(12) = C(15), and C(16) = O(20) groups, which clarifies their great potential for coordination with metal ions. The charges on coordinating sites in the ligand-keto form were O(13)(-0.397459), O(14)(-0.404351), N(12)(-0.162038), and O(20)(-0.372683) (Al-Qahtani et al., 2021). While, in its enol form, the values were O(13) (-0.412866), O(14) (-0.553187), N(12) (-0.352544), and O(20) (-0.352544). The values in the two forms are near enough to allow for similar coordination opportunities. In the Cu(II) complex, Mulliken charges on the coordinating sites are O(13)(-0.47859), O(14)(-0.376319), N(12)(-0.328716), and O(20)(-0.415370), while in the Hg(II) complex, the charges are O(13)(-0.371977), O(14)(-0.240669), N(12)(-0.272094), and O(20)(-0.238556). Except for N(12) and O(20) atoms in the Cu(II) complex due to M → L charge transfer, the charges were essentially lowered from those in the original keto-form due to their coordination towards the metal ions. In the keto form, the groups C(8) = O(13), C(10) = O(14), N(12) = C(15), and C(16) = O(20) have bond lengths of (1.207883), (1.209620), (1.276331), and (1.228438), respectively. As known, the normal property noticed following the coordination of groups, is the elongation of their bonds which already observed in the two complexes.

Furthermore, the formation energy (E, a.u.), dipole moment (D, Debye), E_HOMO (eV), E_LUMO (eV), and ΔE (E_LUMO -E_HOMO, eV) for both tautomer forms were calculated to be −1100.56; 9.294; −0.2198, −0.0981 and −0.123 in the keto form, while, −1100.49, 5.552, −0.2128, −0.1024 and −0.110 in the enol form. The formation energy and the energy gap between frontier orbitals (ΔE) reflect the comparable stability of the two forms (Almalki et al 2021). Also, the keto form's large dipole moment aids the coordination, as happened with Cu(II) and Hg(II) ions. Furthermore, these parameters were recovered for the two complexes, with values of −3197.44, 13.21, −0.1695, −0.1402 and −0.029 in the Cu(II) complex and −1790.58, 13.39, −0.1669, −0.1463 and −0.021 in the Hg(II) complex, respectively. The Cu(II) complex's high dipole moment value underlines its high polarity, expects its lower lipophilicity or miscibility in living-cell lipids. As a result, the biological efficiency of the Cu(II) complex may be ineffective, and the Hg(II) complex, regardless of its features, its efficiency in biological filed may be less important.

3.8 Cyclic voltammetry

The equations used for this study were presented in the supporting materials (part S1), to minimize the similarity percentage in the document.

3.8.1

3.8.1 Cyclic voltammetry of copper chloride

3.8.1.1

3.8.1.1 Redox mechanism and effect of different Cu(II) concentrations without H₃L at 298.15 K

As shown in Fig. S10, the redox mechanism of CuCl₂ in 0.1 M KCl as a supporting electrolyte and glassy carbon (GC) as a working electrode was investigated using cyclic voltammetry within the potential window of 0.8 to −0.8 V and scan rate of 0.1 V/sec at 298.15 K. The emergence of two cathodic and two anodic peaks correspond to (Cu²⁺/Cu¹⁺) and (Cu¹⁺/Cu⁰) on the voltammogram indicates that the Cu(II) solution is electroactive. The two cathodic potential peaks of Cu²⁺/Cu¹⁺ (E_pc1) and Cu¹⁺/Cu⁰ (E_pc2), were appeared at 0.101 V and −0.322 V in the forward scan, respectively. While, the two anodic potential peaks of Cu⁰/Cu¹⁺(E_pa2) and Cu¹⁺/Cu²⁺(E_pa1)were appeared at 0.035 V and 0.212 V in the reverse scan, respectively. As a result, the reaction's redox mechanism can be expressed as follows:

Reduction:
Cu2+ + e-		Cu1+ at 0.101 V (1)
Cu1+ + e-		Cu at −0.322 V (2)
Oxidation:
Cu		Cu1+ + e- at 0.035 V (3)
Cu1+		Cu2+ + e- at 0.212 V (4)

Tables S2 and S3 show the estimated and collected thermodynamic parameters for both anodic and cathodic peaks (A and B). With increasing CuCl₂ concentration, the majority of the examined parameters I_p, Γ, E° and Q for anodic and cathodic peaks were elevated. The existence of the diffusion process is shown by the increase in all peak currents (Shaikh et al., 2006).

3.8.1.2

3.8.1.2 Cyclic voltammetry response of Cu(II) in presence of H₃L at 298.15 K

As shown in Fig. S11, we investigated the electrochemical behavior of the complexation between H₃L and CuCl₂ in 0.1 M KCl solution within the + 0.8 to −0.8 V potential range. By increasing the ligand concentration and favoring the covalence between the copper ions and H₃L, all of the current peaks for the generated redox peak are reduced. The peak potentials for both cathodic and anodic peaks have also been altered. These outputs assumed that the copper and H₃L have a strong affinity to form a stable complex. Additionally, Tables 3 and S4 summarize the values of peak potential of cathodic (E_Pc) and anodic (E_Pa), anodic peak potential separation (ΔE_p), peak currents (i_Pa, i_Pc), peak current ratio (i_Pa/i_Pc), standard potential (E˚), standard heterogeneous electron transfer rate constant (k_s), charge transfer coefficient (α_na), surface coverage (Γ), and the quantity of charge (Q) for CuCl₂ that obtained from previous voltammogram under influence of H₃L in different concentrations.

Table 3 The parameters of Kinetic and Thermodynamic for CuCl₂ wave (A) in presence of H₃L and at 298.15 K and 0.1 scan rate.

Mx106 (mol/ cm3)	Lx106 (mol/cm3)	Epa (volt)	Epc (volt)	ΔEp (volt)	-Ipa x105 (Amp)	Ipc x105 (Amp)	Ip,a/Ipc (Amp)	E˚	Da x106 (cm2/s)	Dc x106 (cm2/s)	αna	Ks x102 (cm/sec)	Γc x109 (mol/cm2)	+ Qc x105	Γa x109 (mol/cm2)	- Qa x105
3.21	0.64	0.194	0.078	0.116	1.66	3.35	0.494	0.136	3.75	0.15	0.782	4.60	11.37	3.45	5.615	1.70
3.18	1.27	0.209	0.067	0.142	1.47	3.22	0.456	0.138	2.98	0.14	0.648	5.23	10.91	3.31	4.978	1.51
3.16	1.90	0.251	0.073	0.178	1.27	2.72	0.468	0.162	2.28	0.10	0.613	6.12	9.222	2.79	4.320	1.31
3.14	2.52	0.250	0.083	0.166	1.26	2.91	0.434	0.166	2.27	0.12	0.650	6.07	9.872	2.99	4.287	1.30
3.13	3.13	0.286	0.091	0.195	0.82	1.66	0.494	0.189	0.97	3.98	0.764	5.00	5.643	1.71	2.787	0.84
3.08	4.62	0.333	0.099	0.235	0.89	1.28	0.701	0.216	1.19	2.42	0.708	5.52	4.331	1.31	3.034	0.92
3.03	6.06	0.328	0.121	0.207	0.97	1.02	0.956	0.224	1.44	1.58	0.738	3.49	3.445	1.04	3.293	0.99
2.94	8.82	0.329	0.134	0.329	1.19	0.85	1.398	0.164	2.30	1.18	0.234	5.60	2.889	0.88	4.040	1.22

3.8.1.3

3.8.1.3 Effect of scan rates on Cu(II) in absence or presence of H₃L at 298.15 K

Figs. S11 and 5 show cyclic voltammograms of Cu(II) in absence or presence of H₃L at 298.15 K in 0.1 M KCl and at various scan rates of 10, 20, 50, and 100 mV/s. It has been discovered that, (i) increasing the scan rate, either cathodic or anodic peak currents rise in a linear manner. (ii) As the scan rate increased, the cathodic peak changed slightly towards the negative potential and the anodic peak shifted slightly towards the positive potential. The linear relationship between the square root of scan rate and peak currents, as seen in Figs. S12 and 6, supports the idea of diffusion-controlled redox process. The ratio of current peak is>1, implying that the system is quasi-reversible (Zaky et al 2012). The solvation and kinetic characteristics for the cathodic and anodic of CuCl₂ in the absence or presence of H₃L at different scan rates are explained in Tables S5, S6, 4 and 5.

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Fig. 5

Cyclic voltammograms of CuCl₂ at different scan rates in absence (A) and in presence (B) of H₃L and at 298.15 K.

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Fig. 6

Relation (i_p vs. ν^1/2) for CuCl₂ in presence of H₃L by the molar ratio (1:1) and different scan rates.

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Table 4 Effect of different scan rates on CuCl₂ [3.23x10^-6 M] in absence of H₃L at 298.15 K for wave (B).

ʋ (V/sec)	Epa (volt)	Epc (volt)	ΔEp (volt)	-Ipa x105 (Amp)	Ipc x105 (Amp)	Ipa/Ipc (Amp)	E˚	Da x105 (cm2/s)	Dcx105 (cm2/s)	αna	Ks x102 (cm/sec)	Γc x108 (mol/cm2)	+ Qc x105	Γa x108 (mol/cm2)	- Qa x105
0.1	0.037	−0.256	0.292	17.7	4.01	4.411	−0.109	4.24	2.177	2.169	0.51	1.362	4.13	6.007	18.2
0.05	0.035	−0.236	0.267	18.1	3.54	5.117	−0.098	8.86	3.383	3.407	0.44	2.401	7.27	12.29	37.2
0.02	0.023	−0.232	0.213	9.26	1.51	6.151	−0.129	5.79	1.530	3.66	0.12	2.554	7.74	15.70	47.6
0.01	0.021	−0.223	0.244	7.81	1.16	6.711	−0.101	8.24	1.829	3.976	0.13	3.948	12.0	26.50	80.3

Table 5 Effect of different scan rates on Cu- H₃L complex for wave (B) and at 298.15 K.

ʋ (V/sec)	Epa (volt)	Epc (volt)	ΔEp (volt)	-Ipa x105 (Amp)	Ipc x105 (Amp)	Ipa/Ipc (Amp)	E˚	Da x105 (cm2/s)	Dcx105 (cm2/s)	αna	Ks x102 (cm/sec)	Γc x108 (mol/cm2)	+ Qc x105	Γa x108 (mol/cm2)	- Qa x105
0.1	0.084	−0.339	0.423	7.33	3.53	2.080	−0.128	7.738	1.789	7.951	3.15	1.196	3.62	2.488	7.54
0.05	0.059	−0.394	0.453	5.57	1.93	2.878	−0.168	8.916	1.076	0.722	0.70	1.312	3.97	3.776	11.4
0.02	0.032	−0.367	0.399	3.18	1.40	2.278	−0.168	7.273	1.402	0.823	0.32	2.367	7.17	5.392	16.3
0.01	0.012	−0.339	0.351	2.83	0.92	3.072	−0.164	11.49	1.217	1.192	0.16	3.120	9.45	9.584	29.0

3.8.1.4

3.8.1.4 Gibbs free energies and the stability constant of Cu- H₃L complex at 298.15 K

By raising the ratio (j) and ligand concentration, the Gibbs free energies (ΔG) and stability constants (log β_MX) for the copper ions that reacted with H₃L to form complex are raised. This indicating greater complexation was occurred as shown in Tables S7 and S8. By adjusting the scan rate, we were able to see how the stability constant changed. As seen in Tables S9 and S10, when the scan rate is reduced, the complex's stability constant increases, indicating a diffusion-controlled response.

3.9 Biological investigation