A series of nickel(II) complexes derived from hydrazide derivatives, electrochemical, thermal and spectral studies

2.1 Synthesis of the organic Ligands

2.1.2

2.1.2 Synthesis of hydrazide derivatives

The investigated hydrazides (Figure 1) were prepared by condensing 1:1 M ratio of acetylacetone, benzaldehyde, cinnamaldehyde, acetophenone, fufuraldehyde, crotonaldehyde or salicylaldehyde, in ethanol with an ethanolic solution of (E)-2-cyano-N-ethylideneracetohydrazide. The reaction mixtures were refluxed on a water bath for 2–5 h in the presence of few drops of glacial acetic acid. The formed precipitates were separated by filtration, washed with 5–10 ml of absolute ethanol and recrystallized from ethanol and dried. The proposed formula of the ligands is based on an agreement between the analytical and mass spectral (Table 1) data as well as confirmed by IR spectra. The ¹HNMR spectra of HBCH, HMCH and H₂SCH in d₆-DMSO showed signals at δ = 9.68–8.70 (S, 1H) for NH proton; signals in the range of δ = 3.98–4.34 ppm for methylene and CH protons of HBCH, H₂SCH and HMCH; and signals in the range of δ = 7.90–6.93 ppm (multiple signals) for aromatic protons. Moreover, the spectrum of H₂SCH showed the OH proton signal at 11.10 ppm (West et al., 1995; Mostafa et al., 2000).

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Figure 1

Chemical formulae of hydrazones.

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Table 1 Abbreviations, full names, melting points, elemental analyses and formula weights (F. W.) of the ligands.

Abbreviation	Full name	Color	m.p. °C	Found (Calcd)(%)		F. W
				C	H	Founda	Calcd.
HCNH	(E)-N-((E)-but-2-enylidene)-2-cyanoacetohydrazide	Yellowish white	135	56.2 (55.6)	5.7 (6.0)	–	151.2
H2CAH	(E)-2-cyano-N-((Z)-4-hydroxy-2-methylpent-3-enylidene)acetohydrazide	White	170	53.6 (55.4)	6.8 (6.7)	179.0	195.2
HCPhH	(E)-2-cyano-N-(2-phenylpropylidene)acetohydrazide	White	120	65.4 (67.3)	5.2 (5.6)	213.0	214.2
HCMH	(E)-2-cyano-N-((E)-3-phenylallylidene)acetohydrazide	Yellow	168	68.0 (67.6)	4.8 (5.2)	213.0	213.2
HCBH	(E)-N-benzylidene-2-cyanoacetohydrazide	White	155	63.3 (64.1)	4.8 (4.8)	188.0	187.2
HCFH	(E)-2-cyano-N-(furan-2-ylmethylene)acetohydrazide	Light yellow	160	54.8 (54.2)	3.6 (3.9)	–	177.2
H2CSH	(E)-N-(2-hydroxybenzylidene)-2-cyanoacetohydrazide	Yellow	190	59.3 (59.1)	4.9 (4.5)	–	203.2

a Values obtained from mass spectra.

2.2 Synthesis of Ni(II) complexes

All complexes were prepared by adding stoichiometric quantities from each ligand (3 mmol) dissolved in ethanol and nickel acetate dissolved (3 mmol) in bidistilled water. The reaction mixtures were refluxed on a water bath for 4–6 h. The precipitate was filtered off, washed with hot ethanol and diethylether, and finally dried in vacuum desiccators over anhydrous CaCl₂.

2.3 Procedure for the pH-metric titration

The titration was carried out at 25 ± 1 °C. The experimental procedure involves the pH-metric titration of the following solutions against the standardized carbonate free sodium hydroxide solution (8.834 × 10⁻³ mol l⁻¹) and 10% (v/v) ethanol–water solution at constant ionic strength (0.1 mol l⁻¹).

• Solution a: 1.5 ml of HCl (0.012 mol l⁻¹) + 2.5 ml of KCl (1 mol l⁻¹) + 2.5 ml of ethanol.

• Solution b: 1.5 ml of HCl (0.012 mol l⁻¹) + 2.5 ml of KCl (1 mol l⁻¹) + 2.0 ml of ethanol + 0.5 ml of ligand (0.01 mol l⁻¹).

• Solution c: solution b + 0.5 ml of (0.001 mol l⁻¹) Ni(II) ion.

For each mixture, the volume was made up to 25 ml with bidistilled water before the titration. The ${\overline{n}}_{\text{A}}$ values are calculated (Agrarwal and Sarin, 1993; Irving and Rossotti, 1954).

Applying these equations: $${\overline{n}}_{\text{A}}=Y+\frac{(V1-V2)([A]+[B])}{(\mathit{Vo}+V1)\mathit{TL}}$$ where Y is the ionizable proton(s) of the ligand, V₁ and V₂ are the volumes of alkali required to reach the same pH in HCl and in the ligand curves, respectively, V_o is the initial volume of the mixture, TL is the ligand concentration in the initial volume, [A] and [B] are the concentrations of HCl and NaOH, respectively. The $\overline{n}$ and pL values were evaluated by: $$\overline{n}=\frac{(V_3-V_2)([A]+[B])}{(V_0+V_2){\overline{n}}_{\text{A}}{T}_{\text{M}}}$$ $$p\text{L}=\log \frac{1+K_1[{\text{H}}^{+}]}{T_{\text{L}}-\overline{n}·T_{\text{M}}}·\frac{V_3+V_0}{V_0}\text{for monobasic ligand}$$ $$p\text{L}=\log \frac{1+K_1[{\text{H}}^{+}]+K_1{K}_2{[{\text{H}}^{+}]}^2}{T_{\text{L}}-\overline{n}·T_{\text{M}}}·\frac{V_3+V_o}{V_o}\text{for dibasic ligand}$$ where V₃ is the volume of alkali required to reach the desired pH in the complex solution and T_M is the initial concentration of the metal ion.

2.4 Chemical and physical measurements

Carbon, hydrogen and nitrogen contents were determined at the Microanalytical Unit of Cairo University. The nickel content was analyzed complexometrically according to the standard methods (Vogel, 1972). IR spectra were recorded on a Mattson 5000 FTIR Spectrophotometer (4000–400 cm⁻¹) using KBr pellets. The UV–Vis, spectra were determined in the DMSO solvent with concentration (1.0 × 10⁻³ M) for the free ligands and their complexes using a Jenway 6405 Spectrophotometer with 1 cm quartz cell, in the range of 200–800 nm. Molar conductance were measured using a Jenway 4010 conductivity meter for the freshly prepared solutions at 1.0 × 10⁻³ moles in DMSO solvent. Magnetic measurements were carried out on a Sherwood Scientific magnetic balance using the Gouy method. The magnetic moment values were evaluated at room temperature (25 ± 1 °C) using a Johnson Matthey magnetic susceptibility balance. The effective magnetic moments were evaluated by applying: ${\mu }_{\mathit{eff}}=2.828\sqrt{{\chi }_{\text{M}}^{-}\text{T}}$ , where ${\overline{\chi }}_{\text{M}}$ is the molar susceptibility corrected using Pascal’s constants for the diamagnetism of all atoms in the complexes and T is the absolute temperature.¹H NMR spectrum of the organic compounds was recorded on a Varian Gemini 200 MHz spectrometer using DMSO-d₆ as solvent. The electron-impact mass spectra of most free ligands and some of their complexes were checked at 70 eV using an AEI MS 30 Mass spectrometer. Thermogravimetric and its differential analysis (TGA/DTG) were carried out in dynamic nitrogen atmosphere (30 mL/min) with a heating rate of 10 °C/min using a Shimadzu TGA-50H thermal analyzer. The molar conductivities of freshly prepared 1.0 × 10⁻³ mol/cm³ DMSO solutions were measured for the soluble complexes using a Jenway 4010 conductivity meter. The protonation constants of the ligands and the formation constants of their complexes at 298 K were determined pH-metrically using a Hanna Instrument 8519 digital pH meter by the Irving–Rossotti method (Irving and Rossotti, 1954). The cyclic voltammetry measurements were carried out with a Potentiostat wave generator (Oxford Press) equipped with a Phillips PM 8043 X–Y recorder. The electrode assembly consists of platinum wires of 0.5 diameter as working and counter electrodes and Ag/AgCl as a reference electrode. Tetrabutylammonium tetraflouroborate (TBA⁺BF4⁻) was used as supporting electrolyte.

3.1 IR spectral analysis

The most significant infrared bands (Table 3) of the ligands and their Ni(II) complexes provide a conclusive evidence for each coordination mode. The ligands coordinate through two different modes based on the orientation of the coordination sites to each others. A mononegative bidentate mode was observed with [Ni(CNH)(OAc)(H₂O)₂], [Ni(CBH)(OAc)(H₂O)₂], [Ni(CPhH)(OAc)(H₂O)₃], [Ni(CMH)(OAc)(H₂O)] and [Ni(CFH)(OAc)]3H₂O complexes. The HCNH, HCBH, HCPhH, HMH and HCFH ligands are coordinating through the C⚌N group and amide oxygen after its enolization process. The lower shift was mainly observed for υ(C⚌N) band by 20–32 cm⁻¹. The disappearance of υ(C⚌O), υ(NH) and δ(NH) bands is observed in the free ligands at: ≈1680, ≈3248 and ≈1480 cm⁻¹, respectively. This disappearance is followed with the appearance of new bands assigned to υ(C–O) and υ(Ni–O) at 1130–1150 and 495–504 cm⁻¹, respectively. New bands observed in 1466–1495 cm⁻¹ and 1323–1400 cm⁻¹ ranges are assigned to υ_as and υ_s (COO⁻) referring to the coordinating acetate. The difference between the two bands is reflecting the acetate coordination mode, in the [Ni(CPhH)(OAc)(H₂O)₃] complex the difference is 177 cm⁻¹ which is referring to its monodentate attachment. All the other complexes reveal a difference between the two υ_as and υ_s bands by 70–96 cm⁻¹ range, which is supporting the bidentate attachment (Abu-Melha and El-Metwaly, 2007). A tridentate mode was observed with [Ni(CSH)(H₂O)] and [Ni(HCAH)(OAc)(H₂O)]H₂O complexes. H₂CSH ligand coordinates as a binegative tridentate mode through C⚌N, O⚌C–NH and OH groups after the enolization of amide and hydroxyl groups. H₂CAH ligand coordinates as a mononegative tridentate mode through C⚌N, O⚌C–NH and OH after the enolization of amide carbonyl. This proposal is supporting with the lower shift of υ(C⚌N) band from ≈1613 to ≈1591 cm⁻¹. The disappearance of υ(C⚌O) band (at ≈1680 cm⁻¹) in the two ligands is followed by the appearance of two bands at ≈1127 and ≈482 cm⁻¹ assigned to υ(C–O) and υ(Ni–O) bands. The lower shift appearance of υ(OH) band from 3480 to 3430 cm⁻¹ as well as δ(OH) from 1360 to 1320 cm⁻¹ is observed in the [Ni(HCAH)(OAc)(H₂O)]H₂O complex. The appearance of new bands in all complexes’ spectra at ≈430 cm⁻¹ is assignable to υ(Ni–N) band. The appearance of two bands in the [Ni(HCAH)(OAC)(H₂O)]H₂O complex spectrum at 1466 and 1390 cm⁻¹ is assigned to the bidentate mode of the acetate group toward the Ni(II). The appearance of bands at 835–858 and 542–580 cm⁻¹ ranges is assigned to ρ_r(H₂O) and ρ_w(H₂O), respectively (Osman et al., 2002) in the investigated spectra. IR data reveal the tendency of all investigated ligands to interact with Ni(II) after the enolization of amide carbonyl which may be expected with the use of acetate salt which slightly raises the basicity of the coordination medium. More or less unshifted υCN band reveals its ruling out from the coordination. This is an expected behavior and may refer to its interaction by resonance with a neighboring amide group which directly is affecting the localization of an unshared pair of electrons sited on its N atom.

3.2 Magnetic and spectral studies

The electronic absorption bands as well as the magnetic moment values are summarized in Table 4. All the structural formulae proposed for the complexes (Figure 2) are based on the data of electronic spectra as well as the magnetic moments which coincide with each other. The ligands have spectral bands in the range of 36,550–26,000 cm⁻¹ corresponding to π → π^∗ and n → π^∗ transitions (Yadav et al., 2003) in DMF solvent. The reflectance spectra were obtained for the complexes in its solid state for assurance on the ruling out of the solvent effect on the complex structure. The electronic spectra of [Ni(CNH)(OAc)(H₂O)₂], [Ni(CBH)(OAc)(H₂O)₂], [Ni(CPhH)(OAc)(H₂O)₃] and [Ni(HCAH)(OAc)(H₂O)]H₂O complexes showed broad absorption bands in the range of 25,220–26,920 and 16,655–17,030 cm⁻¹ that are assigned to ³A₂g → ³T₁g (p)(υ₃) and ³A₂g → ³T₁g (υ₂) transitions, respectively. This indicates the Ni(II) ion complexes in their six-coordination sphere (Bailar et al., 1975). The magnetic moment values of the complexes are in the range of 2.33 to 2.8 B.M. The values are slightly lower than the octahedral range (2.9–3.4 B.M.) which may be due to the absence of orbital–orbital contribution and the value for spin only moment (Aswar et al., 1998). The spectra of [Ni(CFH)(OAc)]3H₂O and [Ni(CSH)(H₂O)] complexes showed absorption bands at 16,775 and 17,230 cm⁻¹ transitions, respectively arising from the ³T₁(F) → ³A₂(F) transitions characteristic of Ni(II) ion in a tetrahedral environment. This agrees with the electronic spectrum of [Ni(Cl)₄]⁻². The tetrahedral geometry is confirmed by the measured magnetic moments (μ_eff = 3.51 and 3.80 B.M.) which are in harmony with the reported (Figgs, 1984) value for the tetrahedral NiCl₂(PPh₃)₂ complex especially with the presence of orbital–orbital contribution. While, the [Ni(CMH)(OAc)(H₂O)] complex spectrum showed an absorption band at 20,100 cm⁻¹ in good agreement with the square–planar geometry of ¹A₁g → ¹A₂g transition this proposal was supported by the diamagnetic appearance of the complex.

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Figure 2

The proposed structures of all isolated complexes. .

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3.3 Mass spectra

The mass spectra of most organic compounds and some investigated complexes have been recorded. The purity of H₂CAH, HCPhH, HCMH and HCBH ligands and [Ni(CNH)(OAc)(H₂O)₂],[Ni(CMH)(OAc)(H₂O)],[Ni(CFH)(OAc)]3H₂O, [Ni(HCAH)(OAc)(H₂O)]H₂O and [Ni(CSH)(H₂O)] complexes were evaluated based on the abstracted data. The spectra of most investigated compounds showed an essential peak attributed to the molecular ion peak confirming that of theoretically proposed. The spectrum of H₂CAH shows a well defined parent peak at m/z = 179.0 (Calcd. 195.2) for (M⁺+1) – H₂O with a moderate intensity. The spectra of HCPhH, HCMH and HCBH organic compounds show a well defined parent peak at m/z = 213 (Calcd. 214.2) attributed to M⁺−1; at m/z = 213.0 (Calcd. 213.2) for molecular ion weight; at m/z = 188 (Calcd. 187.2) for M⁺+1, respectively, with a moderate intensity. The mass spectra of [Ni(CNH)(OAc)(H₂O)₂], [Ni(HCAH)(OAc)(H₂O)]H₂O and[Ni(CSH)(H₂O)] complexes did not display peaks that refer to the molecular ion peak. This may be due to the sudden fragmentation which happened during the evaporation process as an introductory step for scanning. I think a cleavage happened with the water molecules surrounding the Ni(II) ion. The first peak at m/z = 267.7 (Calcd. 303.9); at 312.0 (Calcd. 347.9) and at 259.8 (Calcd. 277.9) for M⁺ – water molecules, with 60.1, 65.3 and 59.3% intensities respectively which is considered higher than that known for a molecular ion peak, usually of a moderate intensity, may support our imagination. The lower intensity gives an idea for the stability of the fragment except the base peak. The molecular ion peak for [Ni(CMH)(OAc)(H₂O)] and [Ni(CFH)(OAc)]3H₂O complexes at m/z = 348 (Calcd. 347.9) and at 347 (Calcd. 347.8) for M⁺ and M⁺-1 respectively, with 23.1 and 28.57% intensities respectively is in a moderate appearance. A representative mass spectrum of the complex [Ni(CFH)(OAc)]3H₂O is shown in Figure 3. The spectrum shows multi peaks corresponding to the successive degradation of the molecule. The molecular ion peak at m/z = 347 (Cacd. 347.8) represents M⁺−1 of the complex. Scheme 1 demonstrates peaks assignable to various fragments arising from the cleavage of the compound. The base peak at m/z = 91.0 (Calcd. 89.8) is assignable to NiCH₃O.

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Figure 3

Mass spectrum of [Ni(CFH)(OAc)]3H₂O.

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Scheme 1

A proposed fragmentation pattern of [Ni(CFH)(OAc)]3H₂O.

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3.4 Thermogravimetric and kinetic studies

The thermogravimetric analysis tool is essential for investigating the probability of attachment of solvent molecules toward the central atom as a crystal or in a coordinate form. The TGA curves of the investigated complexes were obtained over the temperature range of 50–800 °C and exhibited several events. In most investigated complexes, the decomposition stages started at a relatively high temperature till ≈150 °C, which may reflect the high thermal stability of these complexes combined with the absence of hydrated solvent molecules. The first decomposition stage in TG curves is attributed to the expulsion of the coordinated water molecules and in good agreement with the found and the calculated percentage of the weight loss. But, the following stages appeared overlapped with each other and prohibit the exact determination of the initial and the final points of the step, which cause a difference between the found and the calculated percentage. However, the final residue is completely assigned to NiO or NiN sometimes polluted with carbon atoms. The thermal behavior of [Ni(CFH)(OAc)]3H₂O and [Ni(HCAH)(OAc)(H₂O)]H₂O complexes is reflecting the lower thermal stability of them in comparison with others combined with the initial decomposition stage starting at ≈35 °C. The weight loss is corresponding to the removal of the hydrated molecules. The TG curve of the [Ni(HCAH)(OAc)(H₂O)]H₂O complex is taken as a representative example for the decomposition of these complexes. Four degradation stages are observed in the TG curve as shown in Figure 4. The mass loss concerning with the first step (30–80 °C) corresponds to the release of a hydrated water molecule of 5.20 (Calcd. 5.18) mass percentage with an activation energy of 154.6 kJ/mol (first order reaction). The second degradation stage (230–284 °C) is consistent with the elimination of the acetate moiety besides the coordinated water molecule of 22.31 (Calcd. 22.15) mass percentage with a concomitant activation energy of 275.4 kJ/mol. The third decomposition stage (471–597 °C) may be assigned to the removal of an organic part (CH₃COHCH) of 16.10 (Calcd. 16.40) mass percentage. The final degradation stage (597–720 °C) is attributed to the removal of an organic residual part (C₆H₇N₃O) attached with the Ni atom of 34.79 (Calcd. 34.80). The residual mass is in agreement with NiO species of 21.60 (Calcd. 21.47%).

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Figure 4

TG and DTG thermograms of [Ni(CAH)(OAc)]H₂O.

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Non-isothermal calculations were used extensively to evaluate the thermodynamic and kinetic parameters for different thermal decomposition steps in some complexes employing the Coats–Redfern (Coats and Redfern, 1964) and Horowitz–Metzger equations (Horowitz and Metzger, 1963). The results of activation enthalpy (ΔH^# = E − RT); the activation entropy (ΔS^# = 2.303 [log (Zh/KT)]R) and the free energy of activation (ΔG^# = ΔH^# − TΔS^#) are given in Table 5, where Z, K and h are the pre-exponential factor, Boltzmann and Planck constants, respectively (Madhu et al., 1989). The kinetic parameters calculated by the Horowitz–Metzger method revealed no significant difference with those evaluated by the Coats–Redfern method. The activation energies could not be calculated for the overlapping or unsuitable steps. The high activation energy values (E) for some complexes indicate that the removed part is strongly bonded to the Ni(II) ion. Most complexes have negative ΔS^# values suggesting that the activated complex is more ordered than the reactants and the reactions are slower than normal(Mahfouz et al., 2001). The positive value of ΔH^# indicates that dissociation of the complexes is accompanied by absorption of heat and so the process is endothermic. The large positive value of ΔG^# points to the fact that the dissociation process is not spontaneous (El-Asmy et al., 2008; Ma et al., 2012).

3.5 Electrochemical measurements

Electrochemistry of [Ni(CFH)(OAc)]3H₂O, [Ni(CBH)(OAc)(H₂O₂] [Ni(HCAH)(OAc)(H₂O)]H₂O [Ni(CNH)(OAc)(H₂O)₂] and [Ni(CMH)(OAc)(H₂O)] complexes was studied using CV in DMSO (0.1 mol L⁻¹ TBATFB). The complexes showed similar features in the investigated potential range of −1.2 to 1.4 V, and displayed two well-defined electrode couples. The results are summarized in Table 6. A representative voltammogram is shown in Figure 5. The complexes showed two successive one electron processes. The first reduction wave of the complexes is safely assigned to the irreversible couple Ni(II)/Ni(I) with E_1/2 of (−0.49)–(−1.25) V (ΔE_p = 0.27–0.85 V) and represented as follows: Ni^IIL + e⁻ ↔ Ni^IL. The second couple with E_1/2 of (−0.82)–(−1.1) V is assigned to the irreversible electrode couple Ni^II/Ni^III by comparison with analogous Ni(II) complexes (El-Shahawi and Smith, 1994). The ratio ip,c/ip,a > 1 at the sweep rate 20–200 mV confirms the irreversible nature of the electrode couple (Paramanik and Bhatacharya, 1997). The irreversible nature of this couple is also confirmed by the linear dependence of the cathodic peak potential (Ep,c) with sweep rate (log υ) as shown in Figure 6A. The product of the number of electrons involved in the reduction process (n) and the corresponding charge transfer coefficient (α) can be determined from the slope. The dependence of the cathodic peak current (ip,c) of the electrode couple Ni^II/Ni^III on the square route of the sweep rate (υ^1/2) suggests a diffusion-controlled electrochemical process (Figure 6B). The peak–peak potential separation (ΔEp) of the electrode couple Ni^II/Ni^III increased with increasing the scan rate confirming the occurrence of a slow chemical reaction and a limited mass transfer following the electrode process. Thus, the electron transfer process is irreversible and the species that initially formed in the electrode process may also react further to give products that are not reoxidized at the same potential as in the first formed species (Huges and Macero, 1974). The dependence of the voltammetric response of Ni^II/Ni^III on the sweep rate, the depolarizer concentration of the analyte as well as the decrease in i_p,c/υ^1/2 is typical of an ECE (electrochemical reaction coupled between two charge processes) type mechanism in which an irreversible first-order chemical reaction is interposed between two successive one-electron charge transfers.

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Figure 5

Cyclic voltammogram of [Ni(CBH)(OAc)(H₂O)₂] in $\text{DMF}-{\text{TBA}}^{+}{\text{BF}}_4^{-}$ at 100 mV/s vs. Ag/AgCl electrode.

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Figure 6

The dependence of the cathodic peak potential E_p,c (A) and the peak currents i_p,c or i_p,a (B) of the electrode couple Ni^II/Ni^I of the complex [Ni(CBH)(OAc)(H₂O)₂].

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3.6 pH-metric studies

3.7 Molecular modeling

An attempt to gain a better insight into the molecular structure of the ligands and their Ni(II) complexes geometry optimization and conformational analysis have been performed by the use of a MM⁺ force field implementing hyperchem 7.5 (HyperChem, 2002). The drawn modeling structures of the complexes displaying the stable stereo structure include the lowest energy level. The total energy calculated for all complexes is as follows: =44.974 for Ni–CNH, =40.9722 for Ni–CBH, =34.721 for Ni–CPhH, = 47.153 for Ni–CMH, =41.6326 for Ni–CFH, =54.6016 for Ni–HCAH and =45.4202 kcal/mol for Ni–CSH complexes (Figure 7). The total energy content calculated for the investigated complexes support the previous discussion concerning with the pH metric study for the complexes’ stability constants. The total energy content calculated reveal that the Ni(II)–CPhH is the most stable complex isolated. However, the Ni(II)–HCAH complex is the less stable one. This may be due to the essential difference between the two tools as the molecular modeling program is implemented on the isolated solid complexes but the pH metric studies were carried out in solution.

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Figure 7

The modeling structures of the investigated complexes.

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