Synthesis, structural investigations and corrosion inhibition studies on Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) complexes with 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide

Abstract

Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) complexes with 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide (Habph) have been synthesized. The complexes were characterized by different physico-chemical and spectral studies viz. molar conductance, magnetic susceptibility measurements, electronic, IR and NMR spectra. The molecular structures of the ligand Habph and its Mn(II), Ni(II), Cu(II), Zn(II) complexes were further confirmed by single crystal X-ray diffraction technique. The Habph acts as a monobasic tridentate ligand coordinating through pyridyl-N, azomethine-N and enolate-O atoms with metal ions. Magnetic moments and electronic spectral studies suggest a high spin octahedral geometry for all the complexes. The ligand molecule exhibits a Z molecular conformation about the >C=N– bond, whereas the metal complexes show E-configuration in their single crystal structures. The presence of inter- and intra-molecular H-bonding and various C–H⋯π interactions stabilize the molecular structure of the metal complexes. The structure of the Co(II) complex has been satisfactorily modeled by density functional theory (DFT) and time dependent-DFT (TD-DFT) calculations. The results of electrochemical impedance spectroscopy (EIS) and adsorption behavior of the ligand and metal complexes show appreciable corrosion inhibition efficiency for mild steel in a 1 M HCl medium. The metal complexes show a better inhibition effect than the ligand.

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1. Introduction

Metal complexes of aroylhydrazone ligands have remained an attractive area of research for inorganic chemists because of their versatile coordination chemistry and capability to generate varied molecular architecture and geometry.^1,2 Acylhydrazones may undergo keto–enol tautomerization and act as tridentate ligands. Moreover, deprotonation of the –NH group present in the ligands is readily achieved during complexation resulting in the formation of tautomeric anionic species having new coordination properties.³ A lone pair of electrons, present in an sp² hybridized orbital of the nitrogen atom of the azomethine group, is of considerable biological and chemical importance. The tridentate coordination mode of the hydrazone ligands make them suitable as bis-chelating mono-nucleating agents for metal ions, preferring an octahedral coordination geometry.⁴ Such organic ligands can act as hydrogen bonding acceptors–donors and provide recognition sites for π–π stacking interactions to form interesting supra-molecular structures when coordinated to metal ions.^5,6

Acylhydrazone Schiff bases readily coordinate with a wide range of transition metal ions, which exhibit interesting physico-chemical, biological and catalytic properties. Aroylhydrazone complexes of transition metal ions are known to provide useful models for elucidation of the mechanism of enzyme inhibition by hydrazine derivatives and for their pharmacological applications.⁷ The biological activity associated with these compounds is attributed due to the presence of –CONHN=CH– moiety.⁸ The acylhydrazone Schiff base complexes of transition metals have shown high catalytic activities in various chemical reactions such as C–N bond formation using Chan-Lam coupling,⁹ transamidation of carboxamides with amines,¹⁰ epoxidation of olefins¹¹ and polymerization of ethylene.¹² A series of iron(III) complexes containing substituted aroylhydrazone ligands have been used in catalytic epoxidation of olefins with tert-butylhydroperoxide.¹³

The corrosion inhibition study of mild steel in acid media has now become an important area of research in industrial and academic fields. Among the available methods of preventing corrosion, the use of inhibitor is one of the most promising methods. Many N-heterocyclic compounds with polar groups and/or π-electrons are efficient corrosion inhibitors in acidic solutions. Such compounds can adsorb on metal surface and form a bond between the N-electron pair and/or π-electron cloud and metal, thereby reducing corrosion in acidic solutions.^14–16 The ability of Schiff base ligands to form stable complexes closely packed in the coordination sphere of metal ion, introduces another class of compounds for corrosion inhibition. The chelate environment with polyfunctional ligands, might plays a significant role in redox behavior and electro-catalytic reduction reactions. A few recent reports indicate that the metal complexes show greater inhibition efficiency than the free ligands.¹⁷ Interaction of transition metal complexes with mild steel is greatly affected by their standard electrode potentials, their reactivity and the nature of the ligand that could stabilize the metal complexes. The effect of the Schiff base N,N′-bis(salicylaldehyde)-1,3-diaminopropane (Salpr) and its corresponding cobalt complex on the corrosion behavior of steel in 1 M hydrochloric acid solution have been demonstrated.¹⁸

Although a few works on the metal complexes of 2-benzoyl pyridine containing Schiff bases have been reported in literature, the work on metal complexes with 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide is virtually absent. In view of the significant role played by the metal complexes of 2-benzoyl pyridine containing Schiff bases in biological systems^19–21 and their interesting structural properties, we have synthesized Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) complexes with 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide. The complexes are characterized by various physico-chemical and spectral techniques. The molecular structure of ligand and some of the complexes have been determined by X-ray crystallography. The ligand Habph and its metal complexes have also been evaluated for their corrosion inhibition properties.

2. Experimental

2.1 Materials and methods

All analytical grade chemicals were obtained from the commercial sources and used without further purification. Benzoyl pyridine (purchased from Sigma-Aldrich, USA), methyl anthranilate, hydrazine hydrate (SD Fine Chemicals, India) and solvents (Merck Chemicals, India) were used as such. The precursor anthranilic acid hydrazide, (NH₂)C₆H₄CONHNH₂ was prepared by the literature method.²²

2.2 Preparation of 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide (Habph)

A methanolic solution (50 ml) of 2-benzoylpyridine (10 mmol, 1.84 g) was added along with 2 drops of glacial acetic acid to the methanolic solution (30 ml) of anthranilic acid hydrazide (10 mmol, 1.51 g) with stirring. The reaction solution was refluxed continuously for 5 h and then concentrated to half of its initial volume by evaporating the solvent. A pale yellow product was crystallized on cooling the above solution to room temperature. The compound was filtered, washed with ethanol and dried in vacuo. The pure compound was recrystallized from hot methanol. Yield (80%). M.p. 140 °C. Anal. Calc. for C₁₉H₁₆N₄O (316.36): C, 72.13; H, 5.11; N, 17.71. Found: C, 72.01; H, 5.12; N, 17.80%. IR (ν cm⁻¹, KBr): ν(NH₂ + NH) 3442s, 3387s, 3216m; ν(C=O) 1637s; ν(C=N) 1598m; ν(N–N) 990w. ¹H NMR (DMSO-d₆; δ ppm): 14.45 (s, NH); 8.90–6.56 (m, Ar–H); 6.80 (s, NH₂). ¹³C NMR (DMSO-d₆; δ ppm): 176.38 (C=O); 150.11 (C=N); 147.12 (C–NH₂); 140.09–114.19 (aromatic carbons). The single crystal structure of above compound was further confirmed by XRD.

2.3 Synthesis of the metal complexes

The metal(II) complexes of Habph were synthesized by reacting 50 ml of methanolic solution of each metal(II) acetates (10 mmol) in a round bottom flask with 50 ml methanolic solution of the ligand Habph (20 mmol, 6.34 g), separately in 1 : 2 (M : L) molar ratio. For the preparation of Mn(II), Co(II), Ni(II) and Zn(II) complexes, the reaction mixture was refluxed for 4–6 h, and then the resulting solution was evaporated slowly at room temperature to crystallize the product. The Cu(II) complex was formed immediately by stirring the reaction mixture at room temperature. The Mn(II), Ni(II), Cu(II) and Zn(II) complexes were obtained as reddish brown, light green, green and yellow crystals, respectively. The single crystal structures of above complexes were determined by X-ray diffraction technique. However, in spite of all the efforts, Co(II) complex could not be crystallized.

2.3.1 [Mn(abph)₂]. Reddish brown, yield (67%). M.p. 269 °C. μ_eff = 5.81 B.M. Anal. Calc. for C₃₈H₃₀N₈O₂Mn (685.65): Mn, 8.01; C, 66.56; H, 4.42; N, 16.35. Found: Mn, 8.10; C, 66.42; H, 4.45; N, 16.28%. IR (ν cm⁻¹, KBr): ν(NH₂) 3451s, 3378s; ν(C=N) 1583m; ν(N=C–O)⁻ 1612m; ν(C–O)⁻ 1345m; ν(N–N) 1016w; ν(M–O) 456m. UV-vis spectrum (λ_max Nujol, nm): 546, 500.

2.3.2 [Co(abph)₂]. Brown, yield (80%). M.p. 279^{d o}C. μ_eff = 4.83 BM. Anal. Calc. for C₃₈H₃₀N₈O₂Co (689.64): Co, 8.54; C, 66.18; H, 4.39; N, 16.25. Found: Co, 8.51; C, 66.02; H, 4.41; N, 16.19%. IR (ν cm⁻¹, KBr): ν(NH₂) 3449s, 3386s; ν(C=N) 1581m; ν(N=C–O)⁻ 1609m; ν(C–O)⁻ 1337m; ν(N–N) 1028w; ν(M–O) 469m. UV-vis spectrum (λ_max Nujol, nm): 663, 492.

2.3.3 [Ni(abph)₂]. Light green, yield (71%). M.p. 251 °C. μ_eff = 2.89 BM. Anal. Calc. for C₃₈H₃₀N₈O₂Ni (689.40): Ni, 08.51; C, 66.20; H, 4.39; N, 16.26. Found: Ni, 8.46; C, 66.06; H, 4.36; N, 16.17%. IR (ν cm⁻¹, KBr): ν(NH₂) 3444s, 3390s; ν(C=N) 1574m; ν(N=C–O)⁻ 1615m; ν(C–O)⁻ 1351m; ν(N–N) 1020w; ν(M–O) 438m. UV-vis spectrum (λ_max Nujol, nm): 971, 596, 393.

2.3.4 [Cu(abph)₂]. Green, yield (63%). M.p. 238 °C. μ_eff = 1.81 BM. Anal. Calc. for C₃₈H₃₀N₈O₂Cu (694.25): Cu, 9.15; C, 65.74; H, 4.36; N, 16.14. Found: Cu, 9.20; C, 65.56; H, 4.38; N, 16.21%. IR (ν cm⁻¹, KBr): ν(NH₂) 3450s 3381s; ν(C=N) 1583m; ν(N=C–O)⁻ 1612m; ν(C–O)⁻ 1330m; ν(N–N) 1018w; ν(M–O) 461m. UV-vis spectrum (λ_max Nujol, nm): 968, 665.

2.3.5 [Zn(abph)₂]·2H₂O. Yellow, yield (65%). M.p. 258^d °C. Anal. Calc. for C₃₈H₃₄N₈O₄Zn (732.12): Zn, 8.93; C, 62.34; H, 4.69; N, 15.31. Found: Zn, 8.88; C, 62.19; H, 4.65; N, 15.26%. IR (ν cm⁻¹, KBr): ν(NH₂) 3438s 3373s; ν(C=N) 1579m; ν(N=C–O)⁻ 1610m; ν(C–O)⁻ 1329m; ν(N–N) 1025w; ν(M–O) 445m. ¹H NMR (DMSO-d₆; δ ppm): 8.08–6.36 (m, Ar–H); 6.78 (s, NH₂). ¹³C NMR (DMSO-d₆; δ ppm): 165.56 (C–O⁻); 151.95 (C=N); 147.03 (C–NH₂); 138.44–112.78 (aromatic carbons).

2.4 Physico-chemical measurements

The metal contents were analyzed by employing standard literature procedures.²³ Carbon, hydrogen and nitrogen contents were determined on an Exeter Analytical Inc. CHN Analyzer (Model CE-440). The molar conductance of 10⁻³ M solutions of the complexes in DMSO was measured at room temperature on a Eutech Con 510 Conductivity meter. ¹H and ¹³C NMR spectra of the ligand and its Zn(II) complex were recorded in DMSO-d₆ on a JEOL AL-300 FT-NMR multinuclear spectrometer. Chemical shifts were reported in parts per million (ppm) using tetramethylsilane (TMS) as an internal standard. All exchangeable protons were confirmed by addition of D₂O. Infrared spectra were recorded in KBr on a Varian 3100 FT-IR spectrophotometer in the 4000–400 cm⁻¹ region. Electronic spectra of the complexes were recorded on a Shimadzu spectrophotometer, model, Pharmaspec UV-1700 in nujol. Magnetic susceptibility measurements were performed at room temperature on a Faraday balance using Hg[Co(SCN)₄] as the calibrant and corrected for diamagnetism.

2.5 Crystal structure determination

Single crystal X-ray diffraction data of the ligand and its Mn(II), Ni(II), Cu(II) and Zn(II) complexes were obtained at 295(2) K, on a Oxford Diffraction Gemini diffractometer equipped with CrysAlis Pro., using a graphite mono-chromated Mo Kα (λ = 0.71073 Å) radiation source. The structures were solved by direct methods (SHELXL-97) and refined against all data by full matrix least-square on F² using anisotropic displacement parameters for all non-hydrogen atoms. All hydrogen atoms were included in the refinement at geometrically ideal position and refined with a riding model.^24,25 The MERCURY package and ORTEP-3 for Windows program were used for generating structures.^26,27

2.6 Computational studies

All the DFT calculations were performed for the Co(II) complex using Gaussian-09 suit of programs. The complex was treated as an open-shell system using spin unrestricted DFT wave functions (UB3LYP),²⁸ i.e. the Becke three-parameter exchange functional in combination with the LYP correlation functional of Lee, Yang and Parr with 6-31G (d,p) basis set for C, H, N and O atoms²⁹ and effective core potentials basis set LANL2DZ (Los Alamos National Laboratory 2 double zeta)³⁰ for the metal atom. The optimized structure was confirmed to be minima on potential energy surface (PES) by performing harmonic vibration frequency analyses (no imaginary frequency found). No symmetry constraints were applied and only the default convergence criteria were used during the geometry optimizations. Based on the optimized geometry, TD-DFT calculations were performed at the same UB3LYP level to calculate the vertical electronic transition energies.

2.7 Corrosion inhibition measurements

The experiments were performed at 50 ppm concentration after optimization the concentration of studied compounds. The EIS data of mild steel sample of same composition in 1 M HCl obtained in other study¹⁷ have been used as a standard one to compare the inhibition efficiency of all the studied compounds. All electrochemical experiments were performed at 303 ± 1 K in a Gamry electrochemical cell with three electrodes connected to Gamry Instrument Potentiostat/Galvanostat with a Gamry framework system based on ESA400 in a frequency range of 0.01 to 100 000 Hz under potentiodynamic conditions, with amplitude of 10 mV peak-to-peak, using AC signal at E_corr (corrosion potential). Gamry applications include software DC105 for corrosion, EIS300 for EIS measurements and Echem Analyst version 5.50 software packages for data fitting. The solutions of the studied compounds were prepared in 10 : 1 (water : DMSO) ratio mixture to ensure solubility. The experiments were measured after 30 min of immersion in the testing solution (no deaeration, no stirring). The working electrode was prepared from a square sheet of mild steel such that the area exposed to solution was 1 cm². Platinum electrode was used as an auxiliary electrode, and standard calomel electrode (SCE) was used as reference electrode. All potentials were measured vs. SCE. We obtained 5 points per decade during the experiment. The composition of mild steel used for corrosion experiments is wt%) C = 0.17, Mn = 0.46, Si = 0.26, S = 0.017, P = 0.019 and balance Fe. The inhibition efficiency of the inhibitor was calculated from the polarization resistance values using the following equation:where R⁰_p and Rⁱ_p are the polarization resistance in absence and in presence of inhibitor, respectively.

The hydrogen gas volume was measured during the corrosion of mild steel in 1 M HCl solution. For this measurement, a burette was filled with electrolyte solution and was inverted over the mild steel electrode. In order to determine the related activation energies, enthalpy and entropy of activation, the potentiodynamic polarization measurements were studied in the temperature range between 303 and 333 K in 1.0 M HCl in absence and presence of all the synthesized compounds.

The adsorption behavior of studied compounds was experimentally investigated by contact angle measurement of acid solution in absence and presence of 50 ppm of all the studied compounds. Aqueous acid solutions with 50 ppm concentration of the studied inhibitors were prepared and the mild steel samples were then immersed into these solutions for 3 h. Upon removal from the solutions, the samples were dried by means of gently nitrogen flow. Contact angle measurements were performed using the static sessile drop method with a Rame-Hart goniometer (Netcong, USA).

3. Results and discussion

It appears from the analytical data of the complexes that the ligand, 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide reacts with metal(II) acetates in 1 : 2 (M : L) molar ratio to form the metal complexes of general composition [M(abph)₂]. The ligand Habph exists in keto form but undergoes keto-enol tautomerization and deprotonates while reacting with metal(II) acetate. The course of reaction may be represented as:

2Habph + M(CH₃COO)₂·xH₂O → [M(abph)₂] + 2CH₃COOH + xH₂O

where, M = Mn(II), Co(II), Ni(II), Cu(II) and Zn(II).

The metal complexes are colored crystalline solids and melt with decomposition in the temperature range 238–279 °C. They are insoluble in common organic solvents viz. methanol, ethanol, chloroform, diethyl ether, benzene and DMF, but are highly soluble in DMSO. The low molar conductance values of 10⁻³ M solutions of all the complexes in DMSO at room temperature (11.02–19.01 Ω⁻¹ cm² mol⁻¹) indicate that they are non-electrolytes.³¹

3.1 Electronic spectra and magnetic moments

The copper(II) complex shows μ_eff value 1.81 B.M., corresponding to one unpaired electron. This complex exhibits two broad d–d bands at 968 and 665 nm which may be assigned to ²B_1g → ²B_2g and → ²E_g transitions. These transitions indicate a distorted octahedral geometry for the complex.³² The μ_eff value of 2.89 B.M. for the nickel(II) complex corresponds to two unpaired electrons in an octahedral environment. It is further confirmed by its electronic spectral bands observed at 971, 596 and 393 nm. These electronic spectral bands correspond to the transitions ³A_2g(F) → ³T_2g(F), → ³T_1g(F) and → ³T_1g(P).³³

[Co(abph)₂] complex shows μ_eff value 4.83 B.M., is fairly close to those reported for three unpaired electrons in an octahedral environment.³⁴ This complex exhibits two d–d bands at 663 and 492 nm, which may be assigned to ⁴T₁g(F) → ⁴A_2g(F), and → ⁴T_1g(P). These transitions are characteristic of a cobalt(II) complex in a six-coordinate octahedral geometry. Electronic absorption spectra of Mn(II) complex shows two d–d bands of very weak intensity at 546 and 500 nm which may be assigned to ⁶A_1g → ⁴T_1g(G) and ⁶A_1g → ⁴T_2g(G). Above complex also shows the effective magnetic moment value of 5.81 B.M., indicating the presence of five unpaired electrons.

3.2 ¹H and ¹³C NMR spectra

The ¹H NMR spectra of Habph (Fig. S1†) exhibits proton signals due to –NH₂ and >NH as singlet at 6.80 and 14.45 ppm.³⁵ The appearance of >NH proton signal at higher region is due to intra-molecular hydrogen bond between pyridyl-N and amide proton.⁹ In Zn(II) complex, the absence of >NH proton signal suggests that the ligand enolizes during complexation. The signal due to –NH₂ protons of Zn(II) complex appears at the similar position as in the ligand, indicating the non involvement of –NH₂ group in coordination. The disappearance of >NH proton of Habph in its D₂O exchanged ¹H NMR spectra confirms its assignment. The signals for aromatic protons occur nearly at the same position in the complex as in the ligand (Fig. S2†).

¹³C NMR spectra of Habph show a cluster of peaks between 140.09–114.19 ppm corresponding to aromatic carbons of pyridine and phenyl ring (Fig. S3†). In Habph, the signals at 176.38, 150.11 and 147.12 ppm, are attributed to C=O, C=N and C–NH₂ carbons, respectively. The C=O carbon signal disappears in Zn(II) complex due to enolization and a new peak appears at 165.56 ppm assigned as C–O⁻. This suggests that the ligand bonds through a carbonylate-O to metal. The C=N carbon signal shows a down field shift and appears at 151.95 ppm in the Zn(II) complex due to bonding of azomethine-N with metal (Fig. S4†). The presence of C–NH₂ signal at the same position in the complex as in the ligand indicates non-participation of this group in coordination with metal. The signals due to aromatic carbons are also shifted to 138.44–112.78 ppm in Zn(II) complex, suggesting involvement of pyridine-N in bonding.

3.3 IR spectra

The absorption regions for all the complexes are more or less similar due to the similarity in coordination modes of the ligand with the metal centre. The characteristic ν(NH), ν(C=O), ν(C=N) and ν(N–N) bands are observed at 3216, 1637, 1598 and 990 cm⁻¹ in Habph ligand.² The absence of ν(NH) and ν(C=O) bands in all the metal complexes, indicates that the ligand enolizes during complexation and bonding occurs through a carbonilate-O. Appearance of a new ν(C–O⁻) band in all the complexes in the range 1329–1351 cm⁻¹, also confirms the bonding of ligand with metal through C–O⁻ group.³⁶

All the complexes show two ν(C=N) frequencies in the ranges 1609–1615 cm⁻¹ and 1574–1583 cm⁻¹. The first new ν(C=N) appears due to enolization of the ligand and the second ν(C=N) occurs at lower wave number (10–20 cm⁻¹) than the ligand due to coordination of azomethine-N to metal ion. In the metal complexes, ν(N–N) is shifted to higher frequencies by 26–38 cm⁻¹ as compare to the ligand, suggesting coordination of one of the nitrogen atom of >N–N< group with metal. The ν(NH₂) bands appearing at 3442 and 3387 cm⁻¹ in Habph, remain unshifted in its metal complexes suggesting no participation of –NH₂ group in bonding. A non-ligand band observed in the range 438–469 cm⁻¹ in all the metal complexes is tentatively assigned to ν(M–O).

3.4 Crystal structure of Habph ligand

Fig. 1a shows the ORTEP diagram of ligand Habph with atomic numbering scheme. The crystallographic data, structural refinement details are given in Table 1. Some selected bond lengths, bond angles and hydrogen bonding parameters of the ligand are given in Tables 2 and 3. The ligand molecule displays a similar Z molecular conformation about the >C=N– bond. The C(13)–O(1) and C(6)–N(2) display bond distances of 1.231(19) and 1.301(19) Å, as reported for double bonds.¹⁰ The N(2)–N(3) bond distance is 1.367 (18) Å, which is slightly shorter than the single bond distance of 1.411(7) Å, showing some double bond character.³⁷ Due to the presence of intra-molecular N(3)–H(3)⋯N(1) hydrogen bond, Py–C(6)=N(2)–N(3)–C(13)=O(1) skeleton in the ligand is almost planar. The torsion angles N(2)–C(6)–C(7)–C(8) and N(3)–C(13)–C(14)–C(15) are 35.78° and −155.25°, indicate that the terminal phenyl rings are twisted out of the C(6)–N(2)-N(3)–C(13)–O(1) plane. Both the terminal phenyl rings exist with a dihedral angle of 62.72°. The dihedral angle formed between phenyl and pyridyl ring is 53.91°, which suggests that the phenyl ring is tilted from the plane comprising of atoms N(1), C(1), C(2), C(3), C(4) and C(5) due to steric hindrance.

Fig. 1 ORTEP diagrams of (a) Habph with intra-molecular H-bonding (b) [Mn(abph)₂] (c) [Ni(abph)₂] (d) [Cu(abph)₂] (e) [Zn(abph)₂]·2H₂O (f) [Co(abph)₂] (optimized) with ellipsoids of 30% probability.

Table 1 Crystallographic data for Habph and its metal complexes

	Habph	[Mn(abph)2]	[Ni(abph)2]·C2H5OH	[Cu(abph)2]	[Zn(abph)2]·2H2O
a R1 = Σ||Fo| − |Fc||Σ|Fo|.b R2 = [Σw(|Fo^2| − |Fc^2|)^2/Σw|Fo^2|^2]^1/2.
Empirical formula	C19H16N4O	C38H30MnN8O2	C40H36NiN8O3	C38H30CuN8O2	C38H34N8O4Zn
Formula weight	316.36	685.64	735.46	694.24	732.12
Temp, K	293	293	293	293	293
λ (Å)	0.71073	0.71073	0.71073	0.71073	0.71073
Crystal system	Monoclinic	Triclinic	Triclinic	Monoclinic	Triclinic
Space group	P21/c	P1̄	P1̄	P21/n	P1̄
a (Å)	11.6274(4)	10.7689(6)	10.5214(6)	13.2388(7)	10.442(5)
b (Å)	8.6240(3)	11.3142(8)	11.2399(8)	10.7929(11)	11.517(5)
c (Å)	15.6397(7)	15.4981(11)	16.4854(11)	24.2294(14)	15.425(5)
α (°)	90	82.460(6)	79.270(6)	90	100.63
β (°)	91.14	85.509(5)	85.438(5)	103.586(5)	102.96
γ (°)	90	67.300(6)	83.472(5)	90	99.08
V (Å^3)	1567.95(10)	1726.2(2)	1899.6(2)	3365.1(4)	1737.9(13)
Z	4	2	2	4	2
Dcalc (g cm^−3)	1.344	1.3191(2)	1.2858(1)	1.370	1.391
μ (mm^−1)	0.090	0.428	0.559	0.696	0.760
F (000)	664.0	710.0	768.0	1436	752
Crystal size (mm)	0.28 × 0.24 × 0.20	0.32 × 0.30 × 0.28	0.34 × 0.32 × 0.30	0.32 × 0.30 × 0.28	0.32 × 0.28 × 0.24
θ range for data collection (°)	3.20–28.95	3.30 to 29.18	3.18 to 29.09	3.21 to 29.09	3.40–29.04
No. of reflections collected	6695	9333	15 109	16 423	14 181
No. of independent reflections (Rint)	4155(0.0189)	9333(0.0345)	10 199(0.0432)	9024(0.0261)	9283(0.0595)
No. of data/restraints/parameters	4155/0/218	9333/0/442	10 199/3/486	9024/1/451	9283/0/476
Goodness-of-fit on F^2	1.007	1.038	1.025	1.021	0.860
R1, wR2^a^,^b[(I > 2σ(I))]	0.0509, 0.1070	0.0769, 0.2027	0.0858, 0.2099	0.0602, 0.1612	0.0544, 0.0670
R1, wR2^a^,^b (all data)	0.0836, 0.1243	0.1444, 0.2572	0.1529, 0.2649	0.0938, 0.1853	0.1451, 0.0843
Largest difference in peak and hole	0.182, −0.146	0.984, −0.338	1.860, −0.603	1.369, −0.575	0.352, −0.381

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Table 2 Selected bond length (Å) and angle (°) for ligand and metal complexes

Habph
Bond lengths
O(1)–C(13)	1.231(19)	N(1)–C(1)	1.335(2)
N(2)–C(6)	1.301(19)	N(1)–C(5)	1.346(19)
N(2)–N(3)	1.367(18)	C(5)–C(6)	1.492(2)
N(3)–C(13)	1.365(2)	C(6)–C(7)	1.487(2)
N(4)–C(15)	1.372(2)
Bond angles
O(1)–C(13)–N(3)	122.5(15)	N(3)–N(2)–C(6)	118.43(13)
O(1)–C(13)–C(14)	124.34(15)	C(6)–C(5)–N(1)	117.04(14)
N(3)–C(13)–C(14)	113.15(14)	C(1)–N(1)–C(5)	118.80(15)
C(13)–N(3)–N(2)	120.64(14)	N(4)–C(15)–C(14)	121.50(17)

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[Mn(abph)2]
Bond lengths
Mn–O(1)	2.141(3)	O(2)–C(32)	1.281(5)
Mn–O(2)	2.119(3)	N(2)–C(6)	1.289(6)
Mn–N(1)	2.323(4)	N(3)–C(13)	1.325(6)
Mn–N(2)	2.241(4)	N(2)–N(3)	1.369(5)
Mn–N(5)	2.334(3)	N(6)–N(7)	1.381(5)
Mn–N(6)	2.210(3)	C(15)–N(4)	1.353(9)
O(1)–C(13)	1.265(5)	O(34)–N(8)	1.362(6)
Bond angles
O(1)–Mn–O(2)	103.49(13)	Mn–O(1)–C(13)	116.8(3)
O(1)–Mn–N(5)	98.89(13)	Mn–N(1)–C(1)	124.7(4)
O(2)–Mn–N(2)	127.93(13)	N(2)–N(3)–C(13)	109.6(4)
N(1)–Mn–N(2)	69.38(14)	N(1)–C(5)–C(6)	115.9(4)
N(2)–Mn–N(5)	88.31(13)	N(2)–C(6)–C(7)	126.1(4)
O(1)–Mn–N(1)	137.13(13)	O(1)–C(13)–N(3)	124.6(4)
O(1)–Mn–N(6)	131.60(13)	N(3)–N(2)–C(6)	121.0(4)
O(2)–Mn–N(5)	142.13(12)	N(2)–C(6)–C(5)	113.8(4)
N(2)–Mn–N(6)	150.33(14)
Mn–N(2)–N(3)	116.2(2)

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[Ni(abph)2]·C2H6OH
Bond lengths
Ni–O(1)	2.092(4)	O(2)–C(32)	1.271(6)
Ni–O(2)	2.089(3)	N(2)–C(6)	1.295(6)
Ni–N(1)	2.112(4)	N(3)–C(13)	1.348(6)
Ni–N(2)	1.981(4)	N(2)–N(3)	1.365(5)
Ni–N(5)	2.117(4)	N(6)–N(7)	1.369(5)
Ni–N(6)	1.983(4)	C(15)–N(4)	1.355(8)
O(1)–C(13)	1.278(6)	O(34)–N(8)	1.369(9)
Bond angles
O(1)–Ni–O(2)	94.69(14)	Ni–O(1)–C(13)	110.0(3)
O(1)–Ni–N(5)	92.82(16)	Ni–N(1)–C(1)	128.7(4)
O(2)–Ni–N(2)	107.59(14)	N(2)–N(3)–C(13)	108.7(4)
N(1)–Ni–N(2)	78.15(16)	N(1)–C(5)–C(6)	115.6(5)
N(2)–Ni–N(5)	97.64(16)	N(2)–C(6)–C(7)	124.6(5)
O(1)–Ni–N(1)	154.94(16)	O(1)–C(13)–N(3)	125.0(5)
O(1)–Ni–N(6)	105.86(16)	N(3)–N(2)–C(6)	120.9(4)
O(2)–Ni–N(5)	154.70(16)	N(2)–C(6)–C(5)	113.8(4)
N(2)–Ni–N(6)	174.99(16)
Ni–N(2)–N(3)	119.0(3)

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[Cu(abph)2]
Bond lengths
Cu–O(1)	2.181(3)	O(2)–C(32)	1.258(4)
Cu–O(2)	2.063(3)	N(2)–C(6)	1.292(5)
Cu–N(1)	2.233(3)	N(3)–C(13)	1.342(4)
Cu–N(2)	1.986(3)	N(2)–N(3)	1.363(4)
Cu–N(5)	2.108(3)	N(6)–N(7)	1.370(3)
Cu–N(6)	1.953(3)	C(15)–N(4)	1.333(4)
O(1)–C(13)	1.256(4)	O(34)–N(8)	1.338(4)
Bond angles
O(1)–Cu–O(2)	92.82(12)	Cu–O(1)–C(13)	108.9(2)
O(1)–Cu–N(5)	93.54(12)	Cu–N(1)–C(1)	130.1(3)
O(2)–Cu–N(2)	106.55(10)	N(2)–N(3)–C(13)	111.1(3)
N(1)–Cu–N(2)	76.09(12)	N(1)–C(5)–C(6)	114.7(3)
N(2)–Cu–N(5)	97.18(11)	N(2)–C(6)–C(7)	124.1(3)
O(1)–Cu–N(1)	151.83(10)	O(1)–C(13)–N(3)	124.6(3)
O(1)–Cu–N(6)	114.11(11)	N(3)–N(2)–C(6)	119.9(3)
O(2)–Cu–N(5)	156.26(11)	N(2)–C(6)–C(5)	115.8(3)
N(2)–Cu–N(6)	169.29(12)
Cu–N(2)–N(3)	118.9(2)

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[Zn(abph)2]·2H2O
Bond lengths
Zn–O(1)	2.093(2)	O(2)–C(32)	1.290(4)
Zn–O(2)	2.077(3)	N(2)–C(6)	1.289(4)
Zn–N(1)	2.211(3)	N(3)–C(13)	1.337(4)
Zn–N(2)	2.075(3)	N(2)–N(3)	1.376(4)
Zn–N(5)	2.257(3)	N(6)–N(7)	1.369(5)
Zn–N(6)	2.063(3)	C(15)–N(4)	1.346(6)
O(1)–C(13)	1.277(4)	O(34)–N(8)	1.365(5)
Bond angles
O(1)–Zn–O(2)	95.97(9)	Zn–O(1)–C(13)	111.86(19)
O(1)–Zn–N(5)	91.07(9)	Zn–N(1)–C(1)	128.7(3)
O(2)–Zn–N(2)	119.83(9)	N(2)–N(3)–C(13)	109.9(3)
N(1)–Zn–N(2)	75.06(12)	N(1)–C(5)–C(6)	116.0(3)
N(2)–Zn–N(5)	91.13(10)	N(2)–C(6)–C(7)	125.3(3)
O(1)–Zn–N(1)	150.46(10)	O(1)–C(13)–N(3)	125.0(3)
O(1)–Zn–N(6)	112.63(9)	N(3)–N(2)–C(6)	119.9(3)
O(2)–Zn–N(5)	149.04(9)	N(2)–C(6)–C(5)	113.8(3)
N(2)–Zn–N(6)	162.71(12)
Zn–N(2)–N(3)	116.2(2)

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Table 3 Hydrogen bond parameters [Å and °] in Habph and metal complexes

D–H⋯A	D–H	H⋯A	D⋯A	<(DHA)
a #1 = −x, 1 − y, 2 − z, −x, #2 = 1 − y, 1 − z.b #1 = 2 − x, −y, −z, −x, #2 = −1 + x, y, z.c #1 = 1 − x, −y, 1 − z, #2 = −x, 1 − y, −z.
Hydrogen bond parameters [Å and °] in Habph
N(3)–H(3)⋯N(1)	0.86	1.93	2.608(2)	134
N(4)–H(4B)⋯O(1)	0.86	2.08	2.723(2)	131
Hydrogen bond parameters [Å and °] in [Mn(abph)2]^a
N(4)–H(4A)⋯N(3)	0.86	2.08	2.719(10)	130
N(4)–H(4B)⋯N(3)	1.05	1.97	2.662(5)	121
N(8)–H(8A)⋯N(7)	0.86	2.06	2.693(5)	130
N(8)–H(8B)⋯O(1)#1	0.86	2.21	2.924(6)	140
C(11)–H(11)⋯O(2)#2	0.93	2.43	3.256(7)	149
C(19)–H(19)⋯O(1)	0.93	2.42	2.755(8)	101
C(38)–H(38)⋯O(2)	0.93	2.42	2.755(6)	101
Hydrogen bond parameters [Å and °] in [Ni(abph)2]·C2H6OH^b
N(4)–H(4A)⋯N(3)	0.69	2.15	2.705(8)	138
N(4)–H(4B)⋯O(2)#1	0.91	2.11	3.019(7)	179
N(8)–H(8A)⋯N(7)	0.90	2.01	2.655(9)	127
N(8)–H(8B)⋯O(1)#1	0.86	2.21	2.924(6)	140
C(19)–H(19)⋯O(1)	0.93	2.42	2.746(7)	101
C(38)–H(38)⋯O(2)	0.93	2.41	2.752(6)	102
C(39)–H(39)⋯N(8)#2	0.96	2.48	3.011(15)	115
Hydrogen bond parameters [Å and °] in [Zn(abph)2]·2H2O^c
N(4)–H(4A)⋯O(4)#1	0.80	2.27	3.062(6)	171
N(4)–H(4B)⋯N(3)	1.05	1.97	2.662(5)	121
N(8)–H(8A)⋯N(7)	0.96	1.99	2.686(6)	127
N(8)–H(8B)⋯O(1)#2	0.74	2.32	3.054(5)	177
C(19)–H(19)⋯O(1)	0.93	2.38	2.737(4)	102
C(38)–H(38)⋯O(2)	0.93	2.41	2.746(4)	101

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The molecular structure of Habph is stabilized by intra-molecular N(3)–H(3)⋯N(1) and N(4)–H(4B)⋯O(1) hydrogen bonding. The inter-molecular C–H⋯π interactions occur between centroid of pyridyl/phenyl ring and protons of phenyl ring with the contact distances of 2.847, 2.995 and 3.207 Å. The intra-molecular C–H⋯π interaction occurs between centroid of phenyl ring and proton of pyridyl ring with the contact distance of 3.450 Å (Fig. S5†).⁹

3.5 Crystal structures of the [Mn(abph)₂], [Ni(abph)₂]·EtOH, [Cu(abph)₂] and [Zn(abph)₂]·2H₂O

A perspective view of the [Mn(abph)₂], [Ni(abph)₂]·EtOH, [Cu(abph)₂] and [Zn(abph)₂]·2H₂O complexes with labeling scheme is given in Fig. 1b–e. In all the complexes, two molecules of mono-anionic ligand coordinate in an N₄O₂ core via a carbonylate-O, azomethine-N and pyridyl-N atoms to the metal center. The NNO donor sites of the tridentate ligand form five-membered chelate rings of the type CN₂OM and C₂N₂M around the metal center. The C(13)–O(1) bond distance of Mn(II), Ni(II), Cu(II) and Zn(II) complexes 1.265(5), 1.278(6), 1.256(4) and 1.277(4) Å, respectively, are longer than the free ligand, owing to the formation of the M–O bond through enolized C–O⁻ group.⁹ Similarly, N(2)–C(6) bond lengths of Mn(II), Ni(II), Cu(II) and Zn(II) complexes are 1.265(5), 1.278(6), 1.256(4) and 1.277(4) Å, respectively. The M–N(2) (imine nitrogen) bond lengths 2.241(4), 1.981(4), 1.986(3) and 2.075(3) Å, and M–N(1) (pyridyl nitrogen) bond lengths 2.323(4), 2.112(4), 2.233(3) and 2.211(3) Å are observed for Mn(II), Ni(II), Cu(II) and Zn(II), respectively. The shorter M–N(2) bond length as compared to M–N(1) indicates that the azomethine-N coordinates more strongly than the pyridyl-N.³⁸ The carbonyl C–O and imine C–N bond distances in the complexes are intermediate between single and double bond suggesting an extended conjugation in anionic ligand after complexation. The M–O and M–N bond lengths fall in the range reported for similar N,O donor octahedral complexes.³⁹ In Mn(II) complex, the bond angles O(1)–Mn–N(2) 70.13(14)°, N(1)–Mn–N(2) 69.38(14)° and O(1)–Mn–N(1) 137.13(13)° indicate distortion from an ideal octahedral geometry.¹⁷ The Ni(II), Cu(II) and Zn(II) complexes also show similar distortion from the octahedral geometry. In all the complexes torsion angles of −172.22 to −163.19° and −162.63 to 175.53° exhibited by O(1)–C(13)–C(14)–C(15) and O(2)–C(32)–C(33)–C(34) indicate that the oxygen of enolate group is in trans position to the nitrogen of NH₂ group. The torsion angles between O(1)–C(13)–N(3)–N(2), (−7.37 to 7.64)°; N(3)–C(13)–C(14)–C(15), (−6.37 to 17.22)°; and N(2)–N(3)–C(13)–C(14), (−179.96 to 172.12)° indicate that the O(1), N(2) and N(3), C(15) are cis to each other while N(2), C(15) is trans to each other.

The intra- and inter-molecular hydrogen bonds observed in the ligand are absent in its metal complexes. However, a number of new intra- and inter-molecular H-bonding interactions are observed in all the metal complexes (Table 3) which stabilize their molecular structures. The molecular structures of the complexes are also stabilized by various kinds of C–H⋯π interactions. In Mn(II) complex, the intra-molecular C–H⋯π interactions occur between the centroid of five membered chelate/phenyl rings and pyridyl/phenyl protons with the contact distances of 2.698, 2.750 and 3.376 Å (Fig. S6†).⁴⁰ The same intra-molecular C–H⋯π interactions occur with the contact distances of 2.895, 3.029, 3.033, 3.243 and 3.384 Å in Ni(II) complex (Fig. S7†), 2.918, 3.035, 3.302 and 3.382 Å in Cu(II) complex (Fig. S8†), and 2.777, 2.995, 3.350 and 3.391 Å in Zn(II) complex (Fig. S9†). The inter-molecular C–H⋯π interactions occur between the centroid of five membered chelate/phenyl ring and phenyl/amine protons in Mn(II) complex with the contact distances of 3.192, 3.258, 3.282 and 3.368 Å.⁴¹ Such inter-molecular C–H⋯π interactions also occur with the contact distances of 2.580, 3.115 and 3.252 Å in Ni(II) complex, 2.634, 2.778 and 2.982 Å in Cu(II) complex, and 2.759, 3.127 and 3.131 Å in Zn(II) complex.

3.6 DFT optimized structure of Co(II) complex

The density functional theory calculations were carried out in gas phase to optimize the structure of Co(II) complex, using coordinates of the crystal structure of Ni(II) complex from the CIF file. The DFT optimized structure is shown in Fig. 1f. The selected bond lengths and bond angles of the complex are given in Table S1.† The bond distances Co–N(1), Co–N(2), Co–O(1), Co–N(5), Co–N(6) and Co–O(2) are 2.199, 2.097, 2.084, 2.193, 2.097 and 2.088 Å, respectively. The observed bond lengths and bond angles around the metal centre (Table S1†) agree reasonably well with other reported Co(II) complexes of N,O donor ligands in distorted octahedral system.^42,43 The bond lengths N(2)–C(6), N(3)–C(13), C(25)–N(6) and C(32)–N(7) are 1.364, 1.308, 1.364 and 1.308 Å, respectively, which correspond to typical double bond characteristic.

The experimental IR spectral data for the Co(II) complex have been correlated with the DFT calculated data based on peak intensities and peak frequencies (cm⁻¹) (Table S2†). Apart from some minor deviations in theoretical group frequencies from the experimental (3–22 cm⁻¹), the theoretical–experimental agreement is satisfactory. Little deviations are expected due to the negligence of anharmonicity in B3LYP method.⁴⁴ The average error in frequencies calculated by B3LYP method is reported to be of the order of 40–50 cm⁻¹.⁴⁵

In order to get a deeper understanding of the electronic transitions, TD-DFT calculations have been performed for the Co(II) complex. The assignments of the calculated transitions to the experimental bands are based on the criteria of energy and oscillator strength of the calculated transitions. In the description of the electronic transitions, only the main components of the molecular orbitals are taken into consideration. The band assignments are given in Table S3† with their oscillator strengths and energies. The results of TD-DFT calculations on Co(II) complex at the UB3LYP level reveal that the band calculated in the region 374–400 nm is due to mixed ligand → metal (LMCT) and intra-ligand (ILCT) charge transfer transitions. The other low energy absorption bands at 1304, 663 and 492 nm are due to d–d transition with smaller oscillator strength. The orbital analysis of above d–d transition suggests that it originates from d_xy, d_yx, d_xz to d_z², d_x² − y² orbitals, as expected for a distorted octahedral cobalt(II) complex (Fig. S10†).

3.7 Corrosion inhibition study

3.7.1 Electrochemical impedance spectroscopy (EIS) and potentiodynamic polarization (PDP) measurements. EIS is a rapid and convenient method for investigation of protective properties of inhibitors on metals. More reliable results can be obtained by this method, since it does not disturb the double layer at the metal/solution interface.⁴⁶ The impedance plots and proposed equivalent electrical circuit of mild steel in 1 M HCl in absence and presence of studied compounds at 303 K are presented in Fig. 2 and 3. The impedance parameters, solution resistance (R_s), polarization resistance (R_p), charge transfer resistance (R_ct), resistance (R_ads), double layer capacitance (C_dl) and inhibition efficiency (E_EIS%) values are given in Table 4.

Fig. 2 (a) Nyquist plot, (b) Bode-magnitude plot, (c) phase angle plot in absence and presence of 50 ppm concentration of all the studied compounds, (d) Nyquist plot of mild steel in 1 M HCl, (e) Bode magnitude plot of mild steel in 1 M HCl, (f) phase angle plot of mild steel in 1 M HCl.

Fig. 3 Electrochemical equivalent circuit used to fit the impedance spectra.

Table 4 Impedance parameters for mild steel in 1 M HCl in the absence and presence of different inhibitors

Inhibitors	Rs (Ω cm^2)	Rct (Ω cm^2)	N	Y0 (μF cm^−2 sec^n−1)	Cdl (μF cm^−2)	Cads (F cm^−2)	Rads (Ω cm^2)	Rp (Ω cm^2)	EEIS%
—	1.3	17.5	0.821	195.5	56.7	−6.51	−1.3	16.2	—
Habph	1.0	117.3	0.824	122.7	49.6	−1.25	−1.2	116.1	86.0
[Co(abph)2]	0.8	127.4	0.826	110.7	45.1	−0.41	−5.2	122.2	86.7
[Ni(abph)2]	1.2	138.6	0.831	95.3	39.5	−0.24	−4.6	134.0	87.9
[Mn(abph)2]	1.0	143.4	0.832	85.2	35.0	−0.43	−4.9	138.5	88.3
[Cu(abph)2]	1.2	159.2	0.834	72.1	29.6	−0.39	−2.1	157.1	89.7
[Zn(abph)2]·2H2O	1.3	176.5	0.835	55.2	22.1	−0.36	−6.7	169.8	90.5

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The polarization resistance (R_p) which is being used to calculate inhibition efficiency is calculated by using relation.

R_p = R_ct + R_ads (2)

where, R_p is polarization resistance, R_ct is charge transfer resistance and R_ads is resistance due to adsorption of intermediate species.

Nyquist and Bode-phase plots of mild steel in uninhibited and inhibited acid solutions containing 50 ppm concentration of all the studied compounds are presented in Fig. 2a–c. According to Nyquist plots obtained (Fig. 2a), a high-frequency (HF) depressed charge-transfer semicircle was observed (one time constant in Bode plot) followed by a well-defined inductive loop in the low-frequency (LF) regions. The increasing diameter of capacitive loop obtained in 1 M HCl in presence of studied compounds indicates the inhibition of corrosion of mild steel. The high frequency capacitive loop may be attributed to the charge transfer reaction. The presence of low frequency inductive loop may be attributed to the relaxation process obtained by adsorption species like Cl_ads⁻ and H_ads⁺ on the electrode surface.⁴⁷ It is apparent from Table 4 that the impedance of the inhibited system amplifies with increasing the inhibitor concentration and the double layer capacitance (C_dl) values decrease with increasing inhibitor concentration. This decrease in C_dl results from a decrease in local dielectric constant and/or an increase in the thickness of the double layer, suggest that inhibitor molecules inhibit the iron corrosion by adsorption at the metal/acid interface.⁴⁸ The depression in Nyquist semicircles is a feature for solid electrodes and often referred to as frequency dispersion and attributed to the roughness and other inhomogeneities of the solid electrode. In this behavior of solid electrodes, the parallel network: charge transfer resistance-double layer capacitance is established where an inhibitor is present. For the description of a frequency independent phase shift between an applied ac potential and its current response, a constant phase element (CPE) is used which is defined in impedance representation as in eqn (3)

Z_CPE = Y₀⁻¹(jω)⁻ⁿ (3)

where, Y₀ is the CPE constant, ω is the angular frequency (in rad s⁻¹), j² = −1 is the imaginary number and n is a CPE exponent which can be used as a gauge of the heterogeneity or roughness of the surface.⁴⁸ Depending on the value of n, CPE can represent resistance R (n = 0, Y₀ = 1/R), capacitance C (n = 1, Y₀ = C), inductance L (n = −1, Y₀ = 1/L) or Warburg impedance Z_ω = σ(ω)^−1/2(1 − j) where σ is the Warburg coefficient (n = 0.5, ). C_dl values derived from CPE and R_ct parameters according to eqn (4) (ref. 48) are listed in Table 4.

C_dl = (Y₀R_ct¹⁻ⁿ)^1/n (4)

The phase angle at high frequencies provides a general idea of corrosion inhibition performance. The more negative the phase angle, the more capacitive the electrochemical behavior. Charge-transfer resistance increment raises the current tendency to pass through the capacitor in the circuit. Also, increment of the phase angle at relaxation frequency in presence of inhibitors indicates the increase of capacitive response. Such a phenomenon could be attributed to high corrosion inhibition activity of inhibitor. Excellent fit with this model has been obtained for all experimental data. As an example, the Nyquist and Bode-phase plot for free acid solution is presented as Fig. 2d–e. The proposed equivalent electrical circuit for fitting of the experimental data is presented as Fig. 3. The equivalent circuit consists of the double-layer capacitance (CPE) in parallel to the charge-transfer resistance (R_ct), which is in series to the parallel of capacitance (C_ads) and resistance due to adsorption of intermediate species (R_ads).

Corrosion behavior of mild steel has been investigated by potentiodynamic polarization curves in 1.0 M HCl, in absence and presence of 50 ppm concentration of all the studied compounds at temperatures between 303 and 333 K. The plots are given in Fig. 4a–d. The inhibition efficiencies (E_PDP%) have been calculated from current density (i_corr) values through the following equation:

where, i⁰_corr and iⁱ_corr are uninhibited and inhibited corrosion current densities. The obtained data are summarized in Table 5.

Fig. 4 Potentiodynamic polarization curves in 1 M HCl in absence and presence of studied compounds at (a) 303 K (b) 313 K (c) 323 K (d) 333 K.

Table 5 Potentiodynamic polarization parameters of mild steel in absence and presence of different inhibitors at different temperatures

Inhibitors	Temperature (K)	−Ecorr (mV vs. SCE)	icorr (μA cm^−2)	βa (mV dec^−1)	βc  (mV dec^−1)	EPDP%
—	303	469	731	73	127	—
Habph	303	460	110	82	140	84.9
[Co(abph)2]	303	468	104	84	142	85.8
[Ni(abph)2]	303	472	95	83	147	87.0
[Mn(abph)2]	303	480	92	81	148	87.4
[Cu(abph)2]	303	487	84	78	149	88.5
[Zn(abph)2]·2H2O	303	489	76	77	151	89.6
—	313	462	1380	79	117	—
Habph	313	477	250	81	134	81.9
[Co(abph)2]	313	504	236	83	147	82.9
[Ni(abph)2]	313	504	210	84	149	84.8
[Mn(abph)2]	313	492	206	72	142	85.1
[Cu(abph)2]	313	496	190	79	153	86.2
[Zn(abph)2]·2H2O	313	478	170	70	158	87.7
—	323	464	1790	82	156	—
Habph	323	483	381	74	169	78.7
[Co(abph)2]	323	506	344	84	171	80.8
[Ni(abph)2]	323	509	311	88	177	82.6
[Mn(abph)2]	323	511	304	72	169	83.0
[Cu(abph)2]	323	506	280	74	167	84.3
[Zn(abph)2]·2H2O	323	503	275	76	166	84.6
—	333	449	2210	69	149	—
Habph	333	465	542	67	208	75.5
[Co(abph)2]	333	454	492	87	194	77.8
[Ni(abph)2]	333	478	437	66	216	80.2
[Mn(abph)2]	333	481	425	67	210	80.7
[Cu(abph)2]	333	487	394	64	187	82.2
[Zn(abph)2]·2H2O	333	494	388	76	193	82.4

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It is observed that both the cathodic and anodic reactions are suppressed with the addition of studied compounds, which suggest that all the studied compounds reduced anodic dissolution and also retarded the hydrogen evolution reaction. It follows from Table 5 that the values of cathodic Tafel slope (β_c) changes with increasing inhibitor concentration, indicates the influence of the compounds on the kinetics of hydrogen evolution. The shift in the anodic Tafel slope β_a may be due to the chloride ions/or inhibitor molecules adsorbed onto steel surface. According to polarization curves, both the anodic and cathodic current density decrease in presence of inhibitors while E_corr values have not considerably changed (maximum change in E_corr is 21 mV). According to Ferreira and others,^49,50 if the displacement in corrosion potential is more than 85 mV with respect to corrosion potential of the blank solution, the inhibitor can be seen as a cathodic or anodic type. Therefore, all the studied compounds act as mixed type inhibitor at 303 K. The results obtained from potentiodynamic polarization show good agreement with the results obtained from EIS.

3.7.2 Thermodynamic activation parameters. To investigate the mechanism of inhibition and to calculate the activation energies of the corrosion process, polarization measurements are taken at various temperatures in the absence and the presence of 50 ppm concentration of all the studied inhibitors. Corresponding data are given in Table 5. The activation parameters for the corrosion process are calculated from Arrhenius type plot according to the following equation:

and from transition state plot according to the following equation:

where, E_a is the activation energy, R is gas constant, h is Planck's constant, T is absolute temperature, λ is Arrhenius pre-exponential factor, ΔH* is enthalpy of activation, N is Avogadro's number and ΔS* is entropy of activation.

The apparent activation energy of the inhibitors has been calculated by linear regression between log i_corr and 1/T (Fig. 5a); the results are presented in Table 6. Inspection of Table 6 shows that apparent activation energy increases on addition of inhibitors in comparison to the blank solution. The increase in E_a can be interpreted as the physical adsorption. Thus, the studied molecules create a barrier to charge and mass transfer. The higher values of E_a in inhibited solution may also be correlated with the increased thickness of double layer, which enhances the E_a values of the corrosion reaction.⁵¹

Fig. 5 (a) Adsorption isotherm plot for log C_R vs. 1/T, (b) adsorption isotherm plot for log C_R/T vs. 1/T, (c) evolution of H₂ vs. time, t in absence and presence of studied inhibitors (h) and (d) variation of contact angle of electrolytic solution with different studied inhibitors at the mild steel surface.

Table 6 Thermodynamic parameters of activation for mild steel in 1 M HCl in absence and presence of different inhibitors

Name of Inhibitor	Conc. of inhibitor (ppm)	Ea (kJ mol^−1)	ΔH* (kJ mol^−1)	ΔS* (J mol^−1 K^−1)
—	—	29.9	27.1	−99.7
Habph	50	43.3	40.7	−70.7
[Co(abph)2]	50	41.9	39.3	−75.7
[Ni(abph)2]	50	41.3	38.7	−78.4
[Mn(abph)2]	50	41.4	38.8	−78.3
[Cu(abph)2]	50	41.8	39.1	−77.9
[Zn(abph)2]·2H2O	50	44.7	42.0	−69.3

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The relationship between log(i_corr/T) and 1/T are shown in Fig. 5b. Straight lines are obtained with a slope (−ΔH*/2.303R) and an intercept of [log(R/Nh) + (ΔS*/2.303R)], from which the value of ΔH* and ΔS* are calculated and presented in Table 6. The positive sign of enthalpy reflects the endothermic nature of steel dissolution process meaning that dissolution of steel is difficult. On comparing the values of entropy of activation (ΔS*) listed in Table 6, it is clear that entropy of activation increases in presence of the studied inhibitors compared to free acid solution. Such variation is associated with the phenomenon of ordering and disordering of inhibitor molecules on the mild steel surface. The increased entropy of activation in the presence of inhibitor indicates that disorderness is increased on going from reactant to activated complex. The increase in values of entropy by the adsorption of inhibitor molecules on metal surface from the acid solution can be regarded as quasi-substitution between the inhibitor molecules in the aqueous phase and H₂O molecules on electrode surface. In such condition, the adsorption of inhibitor molecules is followed by desorption of H₂O molecules from the electrode surface. Thus increase in entropy of activation is attributed to solvent (H₂O) entropy.

3.7.3 Hydrogen evolution. The volume of hydrogen gas evolved from the reaction system is measured by the gasometric method. This technique provides a rapid and reliable means for assessing the inhibition capabilities of the inhibitor on the mild steel corrosion in an acidic media.⁵² The hydrogen volume as a function of reaction time at 50 ppm concentration of the studied ligand and its complexes is also studied and given in Fig. 5c. From the volume of hydrogen gas evolved, the inhibition efficiency (E_HE%) is calculated by using the equation:⁵³where, V₀ and V_i are the volumes of hydrogen in the uninhibited and inhibited solutions, respectively. From the Fig. 5c, it is indicated that with the addition of studied compounds to the aggressive media reduces the rate of hydrogen gas evolution, indicating that the inhibitor molecules are adsorbed strictly on the mild steel surface and have blocked the electrochemical reaction efficiently through decreasing the available surface area. The results obtained from this technique are in a good agreement with the rest of electrochemical results.

3.7.4 Contact angle of acid solution on mild steel surface. Fig. 5d presents the contact angle of acid solution in absence and presence of studied compounds on mild steel surface. For acid solution without inhibitor, contact angle θ is the lowest (18.9) thereby metal showing most hydrophilic nature. Contact angle decreases in presence of inhibitors. By increasing contact angle, metal shows hydrophobic character to the acid solution containing inhibitor.

4. Conclusion

In the present paper, a Schiff base, 2-amino-benzoic acid (phenyl-pyridin-2-yl-methylene)-hydrazide and its Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) complexes have been synthesized and characterized by various spectral techniques. Molecular structures of most of the compounds have also been determined by X-ray crystallography. The ligand exhibits mono-anionic tridentate behavior and coordinate via a deprotonated carbonylate-O, azomethine-N and pyridyl-N atoms to the metal center. All the metal complexes have a six coordinate distorted octahedral geometry around the metal centre. The structure of Co(II) complex has been optimized by DFT and TD-DFT calculations. The ligand and its metal complexes show appreciable corrosion inhibition property of mild steel in acidic medium.

Acknowledgements

The authors thank the Head, S.A.I.F., Indian Institute of Technology, Kanpur, India for recording the ESI-mass spectra. One of the authors (V.P.S.) is thankful to UGC, New Delhi for financial support and three of the authors (D.P.S., K.T. and M.M.) are thankful to CSIR, New Delhi for awarding Senior Research Fellowship.

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Footnote

† Electronic supplementary information (ESI) available: ¹H & ¹³C NMR spectra of Habph and Zn(II) complex, and π⋯π & C–H⋯π interactions in different compounds; Tables for theoretical structural parameters. CCDC 894885, 988159, 988158, 988160 and 894890. For ESI and crystallographic data in CIF or other electronic format see DOI: 10.1039/c4ra11929k

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