Enhanced Corrosion Protection as a Sustainable Approach for Nickel Using Novel FeL Salen Complex: Electrochemical Investigation and DFT Insights

Abstract

A novel tera-dentate salen ligand and its Fe(III) complex was synthesized and characterized via several spectroscopic and physicochemical techniques. The corrosion rate inhibition of nickel and its alloys inspired the utilization of the L ligand and its FeL complex as vital and eco-friendly inhibitors. To assess their effectiveness, both Tafel plot analysis and electrochemical impedance spectroscopy were employed to examine the electrochemical properties of L and the FeL complex. The results show that corrosion current density (I_corr) steadily drops when the additive concentration is increased, but the inhibition efficiency increases. It has been observed that the efficiency of inhibition rises with temperature, particularly at high temperatures (55 °C) when 1 × 10⁻³ M of L and FeL are present as additives, with η = 90.5% and 92.7%, respectively. Additionally, the findings propose that the adsorption mechanism of both L and FeL additive reptiles follows the Langmuir design isotherm. Electrochemical impedance spectroscopy has also verified these findings. DFT calculations were employed to prove the structure of the investigated FeL complex and its activity as a corrosion inhibitor.

1. Introduction

Essential components of the structural makeup of many materials used in domestic and industrial equipment are metals and their alloys. As these metallic elements come into contact with various harsh conditions, corrosion is a highly important problem in both industry and residential buildings. Numerous direct and indirect costs to industry and humanity are caused by corrosion. One of the most popular ways to stop corrosion is to employ artificial and environmentally friendly corrosion inhibitors in acidic environments. In an acidic solution, many organic compounds with heteroatoms such as N, P, S and X, as well functional groups made of heteroatoms, operate as corrosion inhibitors (CIs). By creating an insulating barrier on the metal, chemical and organic corrosion inhibitors reduce the corrosive action of inorganic acids. For industrial purposes, one of the most crucial materials is nickel and alloys containing nickel. Numerous studies have examined the corrosion of nickel in aqueous electrolytes, driven by the significant technical value of Ni-containing alloys like stainless steel and other alloys faced with corrosion [1]. In many industrial operations, unwanted scale and rust are pickled using acid solutions. Owing to its anti-scaling properties, sulfuric acid is frequently employed for cleansing metals and their alloys in a variety of baths for pickling. Despite the widespread belief that nickel is a “metal that corrodes easily” in harsh environments, it nonetheless corrodes significantly in neutral or mineral acidic environments [2]. Among the best methods to lessen metal corrosion is the use of utilize inhibitors. Interactions between the metal and inhibitor change the properties of the metal’s surface and make it more resistant to severe environments. It is widely recognized that organic matter can prevent the metal from becoming rusted. A great deal of investigation has been done on how different chemicals and inorganic enhancements can impede the action of nickel and its alloys in acidic surroundings [3,4].

Constituents that contain heteroatoms and specific organic compounds, particularly those enriched with phosphorus, nitrogen, sulfur, and oxygen, exhibit remarkable efficacy as corrosion inhibitors [1]. Notably, Schiff bases distinguish themselves in this regard due to the presence of an azomethine group along with sulfur and/or oxygen atoms, all contributing to their potent inhibitory properties, which are further enhanced by the electronegative nature of nitrogen. These compounds work by the particularity of the metal surface’s interaction with the structural categories. Therefore, understanding the interactions that occur between inhibitor molecules and metal surfaces is crucial to developing novel and powerful corrosion inhibitors. Tetradentate dibasic chelating Schiff bases are very good at protecting against corrosion since they have multiple aromatic systems and two imine groups. In addition to this, certain sorts of chemicals have antimicrobial qualities. Because of their low toxicity, they are environmentally friendly inhibitors. These compounds typically attach to metal surfaces and start working [5,6,7,8].

The first stage of an inhibitor’s mechanisms of action in harsh acid exposure environments has been historically speculated about. The phenomenon of an inhibitor’s adsorption on a metal substrate is crucial for understanding how inhibitors operate. This process is governed by various factors such as the molecular charge distribution, the nature of strong electrolytes present, the types of bonds within the organic compound, and the primary interactions that occur at the interface between the organic inhibitors and the metal surface.

Reactive centers such as N atoms and aromatic rings with delocalized p-electron systems can help the pharmaceutical compounds under investigation adsorb onto metal surfaces. These compounds are non-toxic, inexpensive, as well as environmentally benign, and they are utilized in treating hypertension diseases. Additionally, because of their huge molecular weights, they can probably successfully cover a larger portion of the metal’s surface area via adsorption, halting corrosion.

In addition to having an inhibitory effect, substances with heteroatoms such as phosphorus, oxygen, nitrogen, and sulfur also adsorb onto the surface of metals. The nucleophile cores of inhibitor compounds are heteroatoms with unbound electron pairs, or π-electrons, which can easily form bonds. The atoms on the outermost layer of the metal behave electrophilically due to its negative charge, which makes adsorption easier. For example, it has been shown that nitrogen-containing substances exhibit a strong tendency to adsorb onto the negatively charged areas of iron surfaces in acidic conditions; contrarily, compounds with S mainly chemisorb on the anodic regions [9,10].

A significant percentage of inhibitor compounds are hazardous to both human health and their surroundings, despite the fact that some have shown clear anticorrosive potential. Globally, there are ongoing worries regarding safeguarding and the environmental effects of corrosion inhibitors used in different industrial contexts.

Nowadays, researchers are working on alternative, environmentally friendly corrosion inhibitors, such as imine and amino acid combinations, organic compounds [11], and rare earth elements [12].

An example of good inhibitors is amino acids, as they are soluble in aqueous media, safe, biodegradable, and reasonably priced. The body of research on amino acids’ ability to prevent metal corrosion is growing [13,14,15,16,17,18,19]. Among the most plausible candidates is cysteine (HSCH₂CHNH₂COOH), as it possesses three groups: a carboxyl group (-COOH), an amino group (-NH₂), and a thiol group (-SH).

The current study aims to investigate the suppression of nickel corrosion in acidic media using imine ligand and its FeL complex to provide an appropriate mechanism for inhibitions through the use of ac impedance spectroscopy as well potentiodynamic polarization. Furthermore, Ni’s charge–discharge processes have been investigated both with and without inhibitors. Studies on potentiodynamic polarization show that these substances provide good protection in 0.5 M H₂SO₄. SEM and EDAX examinations were used to further investigate the inhibition caused by these chemicals as an occasion of metal surface adsorption. The electrochemical results have been supported by the calculation of specific quantum chemical parameters. DFT insights were applied to confirm the inhibition activity of the compounds under inspection.

2. Experimental

2.1. Reagents

No additional purifying procedures were required because all of the chemicals and solvents utilized in the investigation were of commercial reagent grade (5-bromosalicyldehyde and 3,4 diaminobenzophenone). The initial materials secured from Sigma-Aldrich included metal compounds like iron (III) nitrate nonahydrate, denoted as Fe(NO₃)₃·9H₂O, which signifies an iron salt with three nitrate ions and nine molecules of water. Furthermore, the list of acquired substances features high-quality organic solvents suitable for spectroscopic applications, such as ethanol, glacial acetic acid, and concentrated hydrochloric acid. Sigma-Aldrich provided 99.9% of the alcohol, piperidine, and acetone.

2.2. Synthesis of L Imine Ligand

The condensation procedure, which included the following steps, produced the tetra-dentate L imine ligand (see Scheme 1). Here 5 millimoles (1.06 g) of 3,4-diaminobenzophenone were mixed with 10 millimoles (2.01 g) of 5-bromosalicylaldehyde, which had been previously dissolved in 100 milliliters of ethanol, to form an alcoholic solution. Subsequently, the reaction concoctions underwent reflux for a duration of two hours at a consistent temperature of 85 degrees Celsius, with persistent agitation throughout the process. After that, it was left to cool until a yellow precipitate developed. TLC was used to track the completion of the reaction mixture (Hexane: Ethyl Acetate) (6:4). After filtering and drying over anhydrous CaCl₂, the isolated products were cleaned with ethanol.

¹H NMR (Figure 1 and Figure S2), (δ, ppm), in DMSO-d_6: δ = 6.21 (d, 2H, H-6,6*), 6.98 (d, 2H, H-7,7*), 7.03 (t, 2H, H-1,1*), 7.51 (d, 1H, H-5), 7.56 (t, 1H, H-3,), 7.63 (s, 2H, H-8,8*), 8.08 (s, 2H, H-4,4*), 12.04, 12.70 (s, 2H/-OH); 9.02 and 8.91 ppm (s, 2CH=N).

¹³C NMR for L (Figure S1), δ = 115.8 (C_q), 118.0 (CH), 120.6 (C_q), 123.0 (CH), 124.9 (CH), 128.2 (CH), 129.9 (CH), 130.1 (CH), 132.2 (CH), 133.7 (CH), 135.5 (CH), 137.6 (C_q), 146.4 (C_q), 150.4 (C_q), 165.8 (C_q), 163.7 (CH=N), 187.0 (C=O) ppm.

2.3. Preparation of LFe Imine Complex

A standard procedure was used to obtain the Fe (III) imine complex. When 10 mmol of Fe(NO₃)₃·9H₂O (4.04 g) metal salt was dissolved in an aqueous ethanolic mixture and 10 mmol of L imine ligand (5.78 g) was dissolved in ethanol, a molar ratio of 1:1 was noted. An hour of refluxing at 85 °C was followed by the addition of a piperidine drop. After the solvent evaporated, the precipitated material was filtered out, vacuum-dried, and extensively cleaned with ethanol to eliminate any last unreacted initial ingredients. The complex’s color was bright dark brown about 30 min later (Scheme 1).

We attempted to obtain single crystals of the imine ligand L and its FeL complex using various solvent systems; however, we faced challenges in growing diffraction-quality crystals. Despite multiple trials with different solvent combinations and slow evaporation techniques, the crystals obtained were not suitable for use in X-ray diffraction studies.

2.4. Equipments

All the physical measurements and conditions that were used are recorded in the supporting documentation (Part 1). Individual requirements were typed in the Section 3.

2.5. Stoichiometry and Formation Constants of Complexes in Solution

Job’s method (mole ratio and continuous fluctuation) was used to evaluate the stoichiometric ratio and stability of complexes in solution [20,21,22]. After combining solutions M and L and allowing them to settle, the absorbance was recorded for each solution, and then they were schemed against either ([L]/[M]) or ([L]/[L] + [M]). The following relation is frequently used to obtain the complexes’ formation constants (K_f) via the spectrophotometric process [23]:

(1)${\mathrm{K}}_{\mathrm{f}}={\displaystyle \frac{\mathrm{A}/{\mathrm{A}}_{\mathrm{m}}}{{(1-\mathrm{A}/{\mathrm{A}}_{\mathrm{m}})}^2\mathrm{C}}}$

where [C] is the metal’s molar concentration, A_m is the greatest absorbance for absorbance peak, and [A] is the absorbance values along either side of randomly assigned absorption peaks. Furthermore, the following equation ΔG = −RT ln K_f (at 25 °C) serves to calculate ΔG* values, which represent the free energy changes under standard conditions. This formula aids in understanding the energetic aspects of a chemical reaction and its spontaneity, where K_f, R, and T stand for the formation constant, gas constant, and temperature in Kelvin, respectively.

2.6. Materials and Methods Utilized in Corrosion Inhibition

The materials and procedures used for corrosion are represented in detail in the Supplementary Data [24]. Charge–discharge processes of Ni in 0.5 M H₂SO₄ solution in the presence and absence of different concentrations of inhibitors were conducted utilizing the Galvanostatic charge–discharge technique at applied different current densities (±1, ±3, ±5, ±10, ±15, ±20, ±30 mA·cm⁻²).

2.7. Computational Approach

The geometrical shapes of the molecules of the ligand (L) and its Fe(III) complex were optimized through density functional theory (DFT) calculations. The B3LYP hybrid functional, combined with the 6-311g(d,p) basis set, was employed for the non-metal atoms (C, O, H, Br, and N), while the LANL2DZ basis set was applied to the Fe³⁺ ion. Frontier molecular orbital (FMO) electronic density plots were generated and visualized using GaussView 6.0 software, Wallingford, Connecticut, USA. Computational analyses were conducted using the Gaussian 09W software package, Wallingford, CT, USA [25]. A comprehensive investigation into the electronic properties and potential active sites of the ligand and its Fe (III) complex was conducted. This analysis included molecular electrostatic potential (MEP) surface mapping and the evaluation of quantum chemical reactivity parameters. Moreover, the study investigated the chemical responsiveness and toughness against corrosion of the compounds by analyzing their HOMO and LUMO energy surfaces. To validate these energy levels, the computation of the Density of States (DOS) spectrum was also performed.

3. Results and Discussion

3.1. Physicochemical Properties of the Prepared L Imine Ligand and Its Fe (III) Complex

3.1.1. NMR and IR of the Investigated Imine Ligand and Its FeL the Complex

The identification of compounds was validated by the use of NMR spectroscopy. ¹H NMR spectra (Figure 1) and ¹H NMR (d, ppm) of the imine L ligand in DMSO-d₆.

The nuclear magnetic resonance spectrum of the processed L imine ligand (Figure 1 and Figure S2) is recorded in the Section 2. The nuclear magnetic resonance spectrum of the L imine ligand exhibited singlet signals at δ = 12.04 and 12.74, which are assigned to the two phenolic -OH groups. The signals related to the two azomethine protons (CH=N) were found at δ = 8.91 and 9.02 ppm [26,27]. Furthermore, for aromatic CH, it displayed multiplet signals of eight protons around 6.21–8.08 ppm.

L imine ligand’s ¹³C NMR spectra show a signal at 163.7 ppm for the CH=N and 187.0 ppm for (C=O). Additionally, the signals that show up in the range of 115.8–155.8 ppm are associated with the phenyl carbons (Figure S1).

In KBr pellets, the infrared spectra of L and its complex were discovered to range from 400 to 4000 cm⁻¹. The synthesis of the ligand and the metal complex can be verified by closely comparing the infrared spectra of the two compounds. These ensembles offer vital data concerning the arrangement of the practical groups attached to the metal atoms. The characteristic infrared spectra of the L ligand and its metal compound, along with their respective identifications, are shown in Figure S3. The many bands that result from CH=N and OH reveal information about the structure of the ligand and how it interacts with the metal. The stretching vibration of C=N is the source of the band seen in the L ligand at 1617 cm⁻¹. During complexation, this band moves to a lower frequency (1607). The shift in this band unequivocally demonstrates that azomethine’s nitrogen atoms aid in the creation of metal chelates [28,29].

This is supported by the development of a band at 445 cm⁻¹, which is the stretching vibration of the M-N bond. Bands at 554 cm⁻¹ represent the stretching vibration of M–O [30]. Free OH stretching vibrations are the cause of the 3438 cm⁻¹ band in the ligand spectra. According to the elemental analysis findings, this musical ensemble, frequently referred to as the υ(OH) stretching vibration of crystalline H₂O molecules, exhibits a notable shift to 3422 cm⁻¹. The emergence of a characteristic peak at 886 cm⁻¹ (OH rocking mode) within the synthesized compound suggests the involvement of coordinated water molecules. Upon the process of coordination, the absorption band associated with the metal ion and phenolic oxygen atoms moves towards lower energy levels, reaching 1246 cm⁻¹. This transition implies that these entities are participating in a coordinative bond within the complex. Additionally, the presence of the L imine ligand is discernible through its absorption pattern in the lower frequency range of the spectrum, located at 1276 cm⁻¹, which may be caused by the stretching vibration of the phenolic group (C–O).

3.1.2. Elemental Analyses and Conductivity Measurement Studies

According to Scheme 1, the L mine ligand functions as a tetra-dentate and forms complexes with Fe (III) ions at a metal to ligand ratio of 1:1. Table 1 describes the generated imine ligand along with the data derived from the elemental analysis of its combinations. At room temperature, the resulting complex is pigmented, rigid, and robust, and it possesses a high degree of stability and is not hydrophilic. Using DMF as a solvent, the molar conductance of the prepared FeL complex was measured at room temperature (Table 1). The value of molar conductance at room temperature of the LFe chelate is 10.75 Ω⁻¹ cm² mol⁻¹, which could be attributed to its non-electrolytic character [31].

3.1.3. Electronic Spectrum

Molecular absorption spectroscopy often plays a pivotal role in interpreting the findings of various analytical methods. By scrutinizing the electronic spectra, researchers can deduce the specifics of the metal ion’s ligand environment. The spectra of the ligands and their corresponding complexes were recorded at a temperature of 298 Kelvin, encompassing a wavelength span from 800 to 200 nanometers. Figure 2 displays the maximum absorption wavelength (λ_max) and molar absorptivity (ε_max). There are two distinct bands in the UV–Vis area, which is described as π→π*, n→π*, at about 259,344 nm [32]. These bands are the consequence of a charge transfer process that occurs inside molecules and which takes place in the ligand, and an intra ligand band transition triggered by the activity of the imine. Amazingly, new bands representing the π→π*, n→π*, and d-d transitions were discovered at 259, 299, and 472 nm, respectively [33].

3.1.4. Moment of Magnetism

Magnetic reactivity studies provide information about the geometrical makeup of compounds. Generally speaking, the magnetic moments discovered for complexes can be used to identify the metal ion’s surrounding cooperative geometry (Table 1). As anticipated or octahedral Fe (III) complexes, the magnetic moment of the LFe complex is 5.52 BM, or five unpaired electrons.

3.1.5. Mass Spectra

Mass spectra are an essential tool for defining the complex requirements related to describing metal species structures. The MS of the L imine ligand (Figure 3a) displayed a [M+] at m/z = 578.06 amu that implies their suggested structure [C₂₇H₁₈Br₂N₂O₃; M. wt = 578.25 g/mol]. The produced LFe mass spectra were captured (Figure 3b), and their stoichiometric compositions were compared.

The molecular ion peaks of the [C₂₇H₂₆Br₂FeN₃O₁₁] species are visible in the mass spectra of the Fe (III) chelate at m/z (784.13) (M⁺¹), and they match their suggested molecular weights. The intensity of the ionic reptiles is shown by the base peaks (Figure 3b, Scheme S1). The ultimate dissociation of the Fe (III) chelate was determined by analyzing the other fragmentation ions identified in Scheme S1 at m/z 694.04, 631.85, 565.85, 407.40, 302.86, 211.25, 108.21, 68.14 and 43.78.

3.1.6. Spectrophotometric Determination of the Stoichiometry of the Prepared Complex

Spectrophotometry has been implemented in both the molar ratio view and the continuous variations method to assess the stoichiometry of the complexes. The information gathered signifies that the synthesized compounds have a 1:1 stoichiometry. The continuous variation strategy’s curves demonstrate a 1:1 molar complexation between the metal ions and ligand (Figure 4), with X_ligand = 0.52 being the greatest mole fraction absorption. Plus, the metal ion to ligand ratio of the produced complexes is validated by the data acquired using the molar ratio technique (Figure 4) [34,35].

The correlation between experimental data and the proposed structure of the prepared LFe complex relies on various spectroscopic, analytical, and physicochemical techniques. Elemental analysis (CHN and metal content) confirms the molecular composition, while FT-IR spectroscopy identifies coordination sites through shifts in functional group vibrations, such as -OH, CH=N, M-N, and M-O bonds. UV-Vis spectroscopy provides insight into the electronic transitions, helping to determine the octahedral geometry of the prepared LFe complex. Magnetic susceptibility measurements distinguish between high-spin configurations, further supporting the structural hypothesis. Mass spectrometry (MS) confirms the molecular weight and fragmentation patterns, whereas NMR spectroscopy provides evidence for ligand coordination and symmetry. By analyzing these experimental results collectively, the proposed structure of the metal complex can be rationalized, ensuring consistency with the observed chemical and physical properties.

3.2. Impact of Concentration of the Inhibitor

3.2.1. Tafel Plot

[L] and FeL function as corrosion inhibitors for nickel in the corrosive medium used, and different techniques are employed to examine their impact. The Tafel polarization method is applied to study the rate of corrosion.

Figure 5 displays the shortcomings of several additive concentrations added to 0.5 M H₂SO₄ along with the potentiodynamic polarization curves that are currently in place. Extrapolated potentiodynamic curves were used to obtain the electrochemical polarization parameters for nickel corrosion, such as corrosion potential (E_corr), corrosion current density (I_corr), and cathodic and anodic Tafel constants β_c and β_a. Table 2 displays the inhibition efficiency IE (%) along with further results.

Furthermore, β_c, β_a, and cathodic and anodic Tafel constants were stimulated. Table 2 records the inhibition efficiency (IE (%)) and other outcomes. Figure 6 demonstrated that the corrosion current density decreases as the concentration of inhibitor increases from 1 × 10⁻⁵ to 1 × 10⁻³ M when the solution under examination—0.5 M H₂SO₄—contains additives. This behavior can be explained by the fact that nickel ions produced by corrosion procedures bound by L and FeL created a nickel compound that dissolves in acidic environments. This phenomenon makes sense when we consider that during corrosion processes, nickel ions bound by L create soluble nickel complexes in acidic conditions. The upgraded inhibiting efficiency came about as a result of two processes occurring at the same time that reduce the production of a dissolving nickel complex and nickel corrosion.

At all additive concentrations, the corrosion potential of Ni (E_corr) significantly shifted toward negative potential as the L concentration increased from 1 × 10⁻⁵ to 1 × 10⁻³ M. In the case of utilizing FeL, the corrosion potential is shown to be shifted in the negative direction at all additive concentrations. However, with an increasing FeL concentration, the corrosion potential is shifted in the positive direction (but still more negative than that of pure medium). A closer examination of Table 2 shows that different additive concentrations have an impact on both β_c and β_a values, showing that important anodic and cathodic spots on the metal’s surface were occluded by adsorbent inhibitor particles. As inhibitor concentration is raised, the presence of inhibitors is reduced in important areas where the corrosion process occurred [36].

Additives affect the reaction pathway and mutually change the anodic and cathodic Tafel constants. That result points out that L and FeL react to different alterations in a comparable way. The phenomenon suggests that when ligands and complex molecules concurrently coat the anodic and cathodic sites, the processes of anodic dissolution and cathodic reduction are significantly reduced. Additionally, it was observed that the extent of protection offered by these inhibitors, as well as the external surface coverage, enhanced progressively as the inhibitor concentration in the solution was elevated from 1 × 10⁻⁵ M to 1 × 10⁻³ M. Consequently, the most suitable concentration for protecting metal against corrosion in acidic environments is considered to be 1 × 10⁻³ M for either L or FeL compounds. By examining polarization curves, it became apparent that the rates of both anodic and cathodic reactions diminished with increasing inhibitor concentrations. This observation indicates that, regardless of the specific additive concentration, the presence of these substances leads to a decrease in the rate of the electrochemical processes involved in corrosion. The adding of inhibitors reduces anodic dissolutions and delays hydrogen evolution process [37]. On the metal’s surface, electrochemical reactions are connected to inhibitor adsorption, and it is known that adsorption is dependent on the inhibitors’ chemical structure. The findings, which are displayed in Table 2, prove that the inhibitor-containing systems and the blank system had different E_corr values. At all additive concentrations, it is crucial to note that E_corr values for the inhibitor-containing systems virtually move towards the cathodic regions in comparison to those of the blank. Regarding the E_corr values of L and FeL, notable variations of approximately −180 mV and −116 mV in comparison to the blank E_corr are observed when measured against the SCE standard. These inhibitors can be grouped as cathodic inhibitors since the displacement in E_corr exceeds 85 mV in the negative direction [38,39].

Additionally, the data in Table 2 suggest that the cathodic and anodic Tafel line slopes (βa and βc) are roughly constant and unaffected by inhibitor concentration, indicating that the metal dissolution mechanism is unaffected by the generated inhibitors. The data clearly illustrate that as the concentration of the inhibitor progressively increases, the current densities associated with corrosion tend to diminish. Specifically, at a concentration of 1.0 mM, the L and FeL inhibitors’ I_corr values were found to be 58.3 and 44.67 µA cm⁻², respectively, as stated in Table 2, which is substantially less than the Ni concentration of 616.6 µA cm⁻² in free H₂SO₄. Moreover, as inhibitor concentrations are gradually increased, corrosion currents are reduced.

Furthermore, the information presented in Table 2 indicates a direct relationship between the rise of inhibitor concentrations and the enhancement of surface coverage. As the inhibitor levels increase, a corresponding growth in the proportion of the surface that is covered can be observed, reducing corrosion current densities in the process. The potentiodynamic measurements’ findings demonstrate that the presence of 1.0 mM of L as well FeL results in increased inhibition efficiency (90.5% and 92.7%, respectively). The wider surface coverage results from more inhibitor molecules being adsorbed on the metal surface and these compounds operating as adsorption inhibitors, as evidenced by the increasing inhibition efficiency and decrease in corrosion currents when inhibitors are present [40].

The inhibition efficiency and corrosion current densities of various chemicals are compared, and the results indicate that the dissolution rate of Ni in a 0.5 M H₂SO₄ solution decreases significantly as the concentration of the inhibitor substance increases. The EIS method’s results are reasonably consistent. The data of two methods reveal that the η% of FeL is greater than that of L at the same concentration. The compounds’ molecular size provides the best explanation for this achievement level. The molecular weight of FeL is 748.13 g/mol, which is substantially larger than that of L, which is 578.25 g/mol. This indicates that although both compounds having similar numbers of active sites, FeL exhibits a greater inhibition efficiency (η%) due to its larger size. The Tafel plot analysis of nickel in a 0.5 M sulfuric acid solution, in both the presence and absence of 1 × 10⁻³ M concentrations of L and FeL, is presented in Figure 5a. This outcome highlights the performance and characteristics of the two inhibitors under the specific conditions and a temperature of 25 degrees Celsius. The additive effect of concentration on the corrosion rate of Ni is shown in Figure 5b,c. It is found that growing the FeL level in an acidic solution shifts the corrosion potential positively; in contrast, with L participation, the tendency moves progressively towards negative outcomes as a result of the reduced rates of anodic dissolution as well cathodic hydrogen evolution reactions, which occur with the presence of the additives. All of the corrosion potentials, however, are greater when the additives are included than when they are not.

3.2.2. Electrochemical Impedance Spectroscopy

The verification of potentiodynamic polarization results was aided by examining the EIS of nickel metal in 0.5 M H₂SO₄ with different concentrations of L and FeL. The complicated plane of the impedance plot at E_corr is shown in Figure 7a–c as well Figure 8a–c. Both with and without varying inhibitor concentrations, one capacitive loop can be seen in the impedance spectra acquired with Ni as the anode. The single semicircle shape of the spectrum showed that there was a corrosive reaction occurring at the interface between the electrolyte and the external surfaces. This figure (Figure 7a and Figure 8a) showed that the impedance response of Ni changed significantly when additives in different amounts were added to corrosive media. This can be explained by the fact that both the lowest and greatest inhibitor concentrations cause an increase in substrate impedance. The semicircle looks to deviate from the ideal capacitive behavior because of roughness considerations, surface heterogeneities, and SO₄²⁻ ion adsorption [41]. Conversely, the semicircle steadily improves as the additive concentration increases. When FeL and L are used simultaneously, this phenomenon is seen. According to the Nyquist plot data, the loop size in H₂SO₄ increases with increasing inhibitor concentrations, indicating that the transfer of charges is more complicated [42,43,44,45].

The parameters obtained from the electrochemical impedance analysis for Ni immersed in 0.5 M H₂SO₄ with varying concentrations of L additives (1 × 10⁻⁵, 5 × 10⁻⁵, 1 × 10⁻⁴, 5 × 10⁻⁴, and 1 × 10⁻³ M) at 25 °C are presented in Table 3. The corresponding plots, which include Nyquist diagrams, Bode plots, and phase angle Bode plots, are shown in Figure 7, with panels (a), (b), and (c) displaying these data. The results indicate a clear trend where the charge transfer resistance undergoes a progressive increase with the elevation of L additive concentration. Additionally, it is observed that the double layer capacitance, denoted as C_dl, exhibits a consistent decrease as the concentration of the additive is increased.

Double-layer capacitance results have been calculated with the use of Equation (3). Furthermore, it was found that solution resistance was nearly constant across all additive concentrations [46]. The nickel corrosion rate in a 0.5 M H₂SO₄ environment is found to be inversely related to the concentration of substance L in the solution. As the quantity of L in the solution becomes greater, the corrosion rate of nickel decreases. Furthermore, a similar tendency is observed when FeL is used, as shown in Figure 8a–c. The occurrence of a time constant associated with electrode operation is indicated by Bode plots, as shown in Figure 7b and Figure 8, which display a single-phase maximum at median frequencies. Furthermore, the phase angle rise was more noticeable when FeL was present, suggesting that FeL provides better corrosion protection for Ni than L in acidic environments [47,48]. Furthermore, surface heterogeneity is minimized as a result of additive particle adsorption at particular sites. A different impedance constant of Ni is obtained (Table 3) when the two additions are added in multiple concentrations to 0.5 M H₂SO₄. This behavior confirms previous data obtained from potentiodynamic polarization.

(2)$f (-{\mathrm{Z}}_{\max }^{//})={\displaystyle \frac{1}{2\Pi C_{\mathrm{d} \mathrm{l}}}}{R}_{\mathrm{Ct}}$

3.2.3. Effect of Temperature

As shown in Figure 9a–f, the Tafel plot was used to examine the additives’ effects at different applied temperatures (25–55 °C) in 0.5 M H₂SO₄ in the presence of different concentrations of additives. The effects of temperature on the corrosion’s current density, I_corr, and inhibition efficiency (IE%) in the presence and absence of L and FeL are demonstrated in Table S1, Table 4 and Table 5. It is evident that an increase in I_corr values follows rising temperatures, and as a result, the corrosion rate of Ni in 0.5 M H₂SO₄ is increased in both uninhibited and inhibited media. Moreover, L and FeL inhibition efficiency both increase with rising temperatures on the surface. A rise in temperature has the ability to increase the tendency of adsorption over desorption. This suggests that the amount of surface coverage was here increased. This trend is obviously seen in the case of FeL inhibition efficiency at different temperatures. Moreover, an improved surface coating at elevated temperatures resulted in a minimized roughness of the metal surface. Tafel variables for the Ni metal in 0.5 M H₂SO₄ at various inhibitor concentrations and temperatures are shown in Table 4 and Table 5. Applying the Arrhenius Formula (3), the activation energy is evident; the findings (E_a) of Ni corrosion were calculated both with and without additives at different concentrations [49,50,51,52].

(3)${ \mathrm{logI}}_{\mathrm{corr}}=L\mathrm{ogA} - {\displaystyle \frac{{\mathrm{E}}_{\mathrm{a}}}{2.303 \mathrm{R} \mathrm{T}}}$

Here, T is the absolute temperature, R is the universal gas constant, and A is the Arrhenius pre-exponential variable. For the 0.5 M H₂SO₄ medium, Figure 10 displays lines that are straight and that span log I_corr vs. 1/T of Ni corrosion both with and without multiple additive concentrations. The log I_corr vs. 1/T straight regression coefficients were all close to 0.99 [53]. Activation energy is calculated and listed as well. The findings reveal the presence of inhibitors in 0.5 M H₂SO₄. The solutions’ E_a values rose at all concentrations, and the range of improvement was proportionate to the additive concentrate ion, indicating that the energetic barrier to corrosion increases in the presence of inhibitors. Building a protective layer on the Ni surface slows down the rate of corrosion because progress in the energy barrier is linked to the adsorption of inhibitors. The activation energy reached 45.8 and 50.7 kJ mol⁻¹ for L and FeL, respectively. This pattern in activation energy data supports the findings of the earlier method of additions derived from room temperature Tafel plots. The E_a data are listed in Table S2.

This relationship highlights how the corrosion process develops when the metal is exposed to different thermal conditions without the presence of any inhibitors. The examination of I_corr across a range of temperatures provides insights into the fundamental reaction kinetics and activation energy involved in the dissolution of nickel in the given acidic solution. The study is applied with multiple concentrations to observe the effect of temperature on corrosion dynamics in the absence of any external protective agents.

3.2.4. Thermodynamic Constants and the Adsorption Isotherm for the Corrosion Process

The inhibitors’ abilities to adsorb are modulated by their structure, the type of metal and its charge, the way the charge is distributed inside the molecule, and the type of electrolytes [54,55,56]. Studying the adsorption isotherm in important ways reveals the relationship between the additive and the Ni surface. As L and FeL increase, so do the surface coverage values θ. High inhibitor adsorption on the Ni surface is the explanation for this result (Figure 11). Further to the adsorption of H₂O molecules on the nickel surface, it is commonly thought that the interaction of organic species with the external layer of nickel in an aqueous environment is facilitated by an extra adsorption process involving an organic inhibitor in the aqueous solution, as described by Equation (4) in reference [57].

(4)$\mathrm{Org}{.}_{\left(\mathrm{sol}\right)}+{(\mathrm{H}}_2{\mathrm{O})}_{\mathrm{ads}.}\to \mathrm{Org}{.}_{(\mathrm{ads}.)}+{\mathrm{x}\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{sol}.)}$

A number of isotherms, including Freundlich Frumkin, Langmuir, and Temkin, were tried to fit surface coverage θ values, where x is the percentage or size fraction of H₂O molecules that were used in place of each organic molecule. The findings with the most appropriate fit were derived using the Langmuir adsorption isotherm, the equation of which is as follows (5):

(5)${\displaystyle \frac{{\mathrm{C}}_{\mathrm{inh}}}{\theta }}={ \mathrm{C}}_{\mathrm{inh}}+{\displaystyle \frac{1}{{\mathrm{K}}_{\mathrm{ads}}}}$

where K_ads is the adsorptiion constant, C_inh. The concentration of an additive and the corresponding values of surface coverage (θ) are derived from potendynamic polarization studies. For different concentrations, the adsorption equilibrium constant (K_ads) values are calculated. The determined K_ads for L and FeL are 1.41 × 10⁵ M⁻¹ and 5.9 × 10³ M⁻¹, respectively, at a temperature of 298 K. These high figures indicate that the adsorption of both L and FeL onto the Ni surface is highly effective. It is evident that the adsorption process is significantly more favorable than the desorption process. The relationship between K_ads and the standard free energy change of adsorption is established through the following Formula (6) provided in [56].

Here, K_ads signifies the equilibrium constant for the adsorption–desorption process, which is directly correlated to the standard adsorption free energy (ΔG°). This equation underscores the strong adsorption tendency of L and FeL onto the nickel surface at the specified conditions.

(6)${\mathrm{K}}_{\mathrm{ads}}={\displaystyle \frac{1}{55.5}}\exp (-{\displaystyle \frac{\Delta {\mathrm{G}}_{\mathrm{ads}}^{\mathrm{o}}}{\mathrm{RT}}}).$

Here, K_ads is the equilibrium parameter for adsorption–desorption operations and C_inh. is concentration of organic additive in acidic conditions. The Gibbs free energy of adsorption (ΔG_ads^o) for Ni at 298 K was determined to be −39.30 and −31.40 kJ mol⁻¹ in the presence of L and FeL, respectively. These significantly negative values suggest strong relationships between the inhibitor molecules and the metal surface, indicating favorable interactions [55]. It is well-established that physisorption occurs at ΔG_ads^o values of ≤−20 kJ mol⁻¹, while chemical adsorption, characterized by the formation of coordinate covalent bonds, is represented by values ≥ −40 kJ mol⁻¹ or lower [56]. It should be highlighted, nonetheless, that there are other explanations for organic molecule adsorption besides physical or chemical ones [57]. In this work, ΔG_ads^o values for FeL and L are closer to −40 kJ mol⁻¹. The process of adsorption on the Ni surface involves a mechanism that is primarily chemisorption, yet with some physical adsorption characteristics. It is plausible to suggest that both L and FeL inhibitors are adsorbed through a combination of chemical and physical interactions, with a substantial role played by chemisorption [56]. Table 6 presents the adsorption isotherm parameters, which indicates the efficiency of these additives on Ni in a 0.5 M H₂SO₄ environment at various concentrations. The graphical representation in Figure 12 illustrates the relationship between the Gibbs free energy of adsorption (ΔG_ads) and temperature (T) for Ni metal when exposed to both types of additives.

Standard enthalpy changes ΔH⁰_ads rates, as have been determined. The interaction energies for Ni in H₂SO₄, when treated with significant amounts of L and FeL, were measured at −29.14 and −31.01 kJ/mol, respectively. This outcome indicates that the adsorption process is exothermic. Moreover, the study validated that the effectiveness of inhibition exhibits temperature-dependent variation when using the chemisorption approach (Figure 12). The standard entropy change ΔS⁰_ads values were found to be −520 and −160 kJ/mol for Ni, respectively. The negative value of ΔH⁰_ads revealed an exothermic adsorption path for organic additives [58,59,60,61]. A decrease in solvent energy as well as a decline in water desorption entropy may be associated with a negative ΔS⁰_ads value [62,63,64].

3.2.5. Surface Characterization

Scanning Electron Microscope (SEM)

The inhibiting impact of inhibitors was assessed by assessing scanning electron imaging SEM microscope images of Ni following a deficit soak in 0.5 M H₂SO₄, while the presence of various concentrations of L and FeL is displayed in Figure 13a,b and Figure 14a–c. With a lack of additives, the metal surface is extremely rough and coated by a heavy coating of spongy oxide from NiO and NiSO₄ due to severe corrosion (Figure 13a). The surface is smoother and the quantities of sulfate and nickel oxide particles are reduced when L is present at a high concentration 1 × 10⁻³ (Figure 13b) compared to when it is not, pointing out that the L additive has an inhibiting impact. When different concentrations of FeL are present in the solution (1 × 10⁻⁵, 1 × 10⁻⁴, 1 × 10⁻³ M), the NiO and NiSO₄ particles undergo a significant transformation, effectively disappearing, and the outer layer of the material transitions to an almost transparent and pristine state. New small particles are present on the surface due to the presence of the FeL complex, which is chemisorbed. The amount of these particles is very low, and the surface in this case was apparent. These outcomes indicate how the inhibition efficiency of FeL increases Ni’s ability to withstand corrosion in acidic settings. This is how a low corrosion rate and the positive usage of Ni in acidic atmospheres were attained via green and well-proposed inhibitors [65,66,67,68,69].

Energy-Dispersive X-Ray Spectroscopy (EDX)

The preservative layers implemented on the Ni surface have been analyzed via EDX analysis. The intensity of the Ni signal suggests the formation of protective layers. Table 7 presents the results of the EDAX analysis of Ni immersed in 0.5 M H₂SO₄ with the presence of 1 mM of L and FeL. The findings indicate that in the absence of additives, the corroded Ni surface consists mainly of metal oxides and metal sulfates (as seen in Table S3). When the electrode is exposed to 0.5 M H₂SO₄ with 1 × 10⁻⁵ M of L inhibitor, there is an increased presence of oxygen and sulfur on the Ni surface compared to after exposure to the acid alone. This suggests that the inhibitor has an effect on the film formation and composition. This maintains the preventive impact of ligand particle adsorption on the metal’s surface. This is an important success, since it validates earlier discussions and data gathered through employing the technique of Electrochemical Impedance Spectroscopy (EIS) alongside Tafel plot analysis, after which it was observed that the addition of 1 × 10⁻³ M of the compound FeL to the system resulted in exceptional protection for the nickel (Ni) surface. This remarkable protection effect is attributed to the formation of a protective film of adsorbed inhibitor molecules (FeL) that adhere to the nickel’s surface, thereby significantly enhancing its corrosion resistance. This led to the outstanding level of inhibiting impact. The ratios of oxygen are decreased from 32.20% in the presence of L to 9.66% by mass in the presence of FeL. Additionally, the percentages of sulfur in oxide films were clearly reduced from 8.21% in the case of L to 3.53% by mass in the case of FeL, while percentages of Ni are highly increased, indicating that a pure surface of the metal has been obtained, while oxides and sulfates have been greatly minimized. An iron content also appeared on the derived surface layers when FeL additive was used, confirming the chemisorption of FeL. Figure 15a,b shows an EDAX analysis of oxide coatings formed on the Ni surface after Tafel experiments in 0.5 M sulfuric acid solution containing different concentrations of FeL additive. Figure S4 shows the EDAX analysis of oxide films acquired on the Ni surface in pure medium without additives.

3.2.6. Charge Discharge Studies

Galvanostatic charge–discharge curves for nickel electrode in 0.5 M H₂SO₄ without and with the inclusions applied at multiple current densities (±1, ±3, ±5, ±10, ±15, ±20 as well as ±30 mA·cm⁻²) are shown in Figure 16a–f and Figure 17a–e. The figure illustrates that when the applied current density is raised from ±1 to ±30 mA·cm⁻², the examined electrode’s specific capacitance S.C. and potential difference ΔV generally rise. This shows that, in both the presence and absence of additives, the electrode under consideration exhibits a desirable charge–discharge behavior; in addition, a suitable discharge level was established at varied current densities. Additionally, it has been observed that in both the presence and absence of additives, increasing the applied current density increases the discharge time (t_dis.).

Figure 18a,b shows the galvanostatic charge–discharge curvatures for the Ni electrode in the absence as well as presence of 1 × 10⁻⁵ and 1 × 10⁻³ M of L and FeL at a conducted current of ±30 mA·cm⁻². It was noticed that the Ni metal reached the highest discharging time (13.3 s) in the presence of 1 mM of L and FeL at an applied current density of ±30 mA·cm⁻². Also, the specific capacitance of Ni reached improved values 0.544 and 0.179 mAhr in the presence of 5 × 10⁻⁵ M of L and FeL, respectively, at the applied current density ±30 mA·cm⁻². It is commonly recognized that the energy density of nickel metal is primarily determined by how well the nickel anode material is used in the solution. Nevertheless, in certain situations, the oxidation of nickel may not be suitable for the nickel anode, and in the solutions used here, hydrogen spontaneously evolves on the surface. As a result, increasing concentrations of L and FeL in the solution slows down the nickel’s spontaneous self-discharge, increasing discharge efficiency. Table 8 shows the charge–discharge parameters (t_dis., ΔV, S.C.) for the investigated electrode in 0.5 M H₂SO₄, possessing different concentrations of L and FeL.

The above results show the stability of the investigated electrode. It was detected that the discharging time of Ni reached a higher value in the presence of 1 mM of L and 0.1 mM FeL in 0.5 M H₂SO₄, showing that its impedance behaviors were improved, and thus confirming the obtained impedance data. These experimental data demonstrate that energy efficiency has been enhanced. Finally, the data show that the addition of L and FeL has a notable impact on the developing features of the nickel anode in acidic media. This improvement can be attributed to the fact that the corrosion potential of E_corr has been switched to an excessively negative value when the electrode is in the presence of these additives. Accordingly, the investigated electrode has lower corrosion attack resistance, enhanced discharging time, as well higher specific capacitance (especially in the presence of 5 × 10⁻⁵ M), in the presence of our novel prepared ligand and its FeL complex.

3.3. Computational Details

3.3.1. Optimized Structure

The geometric optimization of L (ligand) and its complex (FeL) was carried out using the DFT/B3LYP method in the gas phase. For this process, the 6-311g (d,p) basis set was employed for lighter atoms, while the LANL2DZ basis set was utilized for Fe³⁺. The calculated lowest-energy configurations for L and FeL were determined to be −6523.40 and −7002.49 Hartree, respectively. The findings reveal that the synthesized Fe(III) complex exhibits significantly enhanced energy stability compared to the L ligand. Furthermore, the optimized structures were confirmed to represent true energy minima, as evidenced by the complete absence of imaginary harmonic frequencies. The optimized geometry of the title compounds is demonstrated in Figure 19. During the optimization process, the Fe³⁺ ion was found to adopt an octahedral geometry. The DFT analysis of the metal ion’s configuration aligns well with the structural model of the FeL complex presented in the Section 2.

3.3.2. MEP Mapping

The MEP offers a visually creative illustration of the electrostatic charge distribution within a molecule, providing valuable insights into the properties of its binding areas [70]. Typically, different colors are used to signify varying charge regions—red, orange, and yellow indicate negatively charged zones prone to electrophilic interactions; blue highlights areas susceptible to nucleophilic attacks; and green denotes regions with neutral potential [71,72]. Using the DFT/B3LYP/6-311g (d,p)/LANL2DZ methodology, MEP diagrams were generated for L and LFe. The results are illustrated in Figure 20. The elected representative from region L exhibits a significant amassing of negatively charged particles, namely, electrons, in the proximity of several atoms, namely, O1, O2, N1, O3, and N2. These atoms display a discernible pattern of charge concentration, with a distinct color gradient transitioning from red to yellow. A region emerges prominently in the area between N1, N2, O1, and O2, representing the presence of lone electron pairs on the electronegative atoms. This observation points to the existence of Lewis base species within these regions. As a result, it is evident that the space surrounding these atoms is highly susceptible to electrophilic interactions [73,74]. Specifically concerning metal ions like Fe³⁺, the Molecular Electrostatic Potential (MEP) map of the LFe system discloses significant information regarding its electrical properties. This visual representation provides a deeper understanding of the complex’s electrostatic nature and interactions. In this map, the O3 atom and the center of the molecule (the region between the O6 and O7 atoms) are highlighted in red-yellow, indicating areas of strong electron density. Conversely, the blue regions correspond to the hydrogens attached to the O4 atom, signifying areas of electron deficiency. Consequently, the electrophilic interactions of the LFe complex are predominantly directed toward the O3, O6, and O7 atoms. The Molecular Electrostatic Potential (MEP) analysis reveals that the nucleophilic propensity is particularly pronounced near the hydrogens attached to the O4 atom. This indicates that the O6 and O7 atoms possess a higher likelihood of engaging with a metallic interface relative to the rest of the molecule. This distribution is advantageous as it enhances the material’s corrosion resistance. The MEP map has been standardized with color scales ranging from −0.84 to +0.84 arbitrary units for L and −1.13 to +1.13 arbitrary units for FeL.

3.3.3. Mulliken Atomic Charge

The Mulliken survey of the population offers insights into the charge distribution across individual atoms within a molecule, making it a valuable tool for examining various aspects of molecular geometry. Accurately determining atomic charges is essential when applying quantum chemical computations to study molecular interactions. For the ligand (L), the Mulliken charges for selected atoms, N1 (−0.342), N2 (−0.444), O1 (−0.347), O2 (−0.409), and O3 (−0.299), were computed based on the geometry optimized in the gas phase. These calculated negative charges on the nitrogen and oxygen atoms indicate a higher electron density in these regions, which makes them more susceptible to electrophilic interactions. As a result, N1, N2, O1, and O2 atoms are likely to form bonds with positively charged species, such as Fe³⁺. This finding aligns with the predictions made by the MEP diagram. The optimized geometries of L and FeL exhibit dipole moment values of 3.12 and 10.49 Debye, respectively. This result indicates that the LFe complex possesses a significantly higher dipole moment, suggesting that the process of complexation induces substantial dipole–dipole interactions within the structural framework of the complex. The Mulliken charge’s dissemination along with the dipole moment direction for the L compound is revealed in Figure 21.

3.3.4. HOMO-LUMO Analysis

Frontier molecular orbitals (FMOs) encompass the most occupied molecular orbital (HOMO) and the least unoccupied molecular orbital (LUMO). These orbitals have significant influence over a molecule’s optical and electrical characteristics, as well as its behavior in quantum chemistry. The HOMO reflects a molecule’s potential to donate electrons, whereas the LUMO signifies its capacity to accept electrons. The HOMO-LUMO energy gap plays a critical role in determining a molecule’s kinetic stability, optical polarizability, chemical reactivity, and chemical hardness or softness [75]. To assess the energetic properties of the title compound, the HOMO–LUMO energy gap was determined using the B3LYP/6-311g (d,p)/LANL2DZ computational method. The computed HOMO and LUMO energy, along with energy gap values of L compound, are obtained as −6.09, −2.92, and 3.17, respectively, in the gaseous phase. Despite this, the related values for HOMO and LUMO energies for the LFe complex are observed at −6.09 and −3.69 eV, respectively. The Fe(III) complex exhibits an energy gap of 2.39 eV, a characteristic that underpins the charge transfer interactions occurring within the molecule. This interaction plays a pivotal role in shaping the metal complex’s chemical reactivity. Figure 22 illustrates the energy gap and surfaces of HOMO and LUMO, as well as their spatial distributions, for the newly synthesized compounds.

3.3.5. Inhibition of Corrosion and Quantum Reactivity

The assessment of a corrosion inhibitor’s characteristics and performance is heavily reliant on a thorough investigation of its electronic configuration and molecular shape. To achieve this, a series of quantum chemical computations were performed on the substances denoted as L and FeL. These compounds were subjected to detailed electronic structure analysis and geometrical optimization to evaluate their corrosion inhibition capabilities. The quantum chemical properties of these compounds were analyzed, focusing on parameters derived from the HOMO and LUMO energy levels. These properties include absolute electronegativity (χ = − (E_HOMO + E_LUMO)/2), absolute softness (σ = 1/η), chemical potential (P_i = − χ), additional electronic charge (ΔN_max = −P_i/η), global electrophilicity (ω = P_i²/2η), and absolute hardness (η = (E_LUMO − E_HOMO)/2) [76,77]. The calculated values for these parameters are summarized in Table 9.

The evaluation of how a substance interacts with a metallic interface can be assessed by considering its E_LUMO and _HOMO (Energy Level of the Lowest Unoccupied Molecular Orbital and Highest Occupied Molecular Orbital) to calculate the energy band gap (ΔE). Typically, an increase in the HOMO energy or a decrease in ∆E correlates with a stronger propensity for the substance to adhere to the surface of the metal [78,79]. According to the data, the LFe complex exhibits a narrower band gap (2.39 eV) compared to the ligand (∆E = 3.07 eV). According to these data, the L ligand complexation, which is indicated by its smaller band gap of 2.39 eV, could enhance its potential to mitigate metal surface corrosion effectively. The global reactivity of the compounds, determined by their energy gaps, reveals a clear trend: LFe exhibits higher reactivity than L. This behavior aligns with the principles of the hard–soft acid–base (HSAB) theory. According to this framework, biological targets such as cells, proteins, and surfaces are classified as “soft” entities. Soft molecules display a stronger affinity for interacting with these targets compared to their harder counterparts [80]. Consequently, increased molecular softness corresponds to enhanced biological activity and heightened chemical reactivity. Therefore, the reactivity hierarchy is predicted to follow the order LFe > L. The absolute electronegativity (χ) parameter serves as an indicator of a compound’s tendency to take up electrons when a back connection is being formed. This dynamic process, involving the interplay of electron donation and back-donation, enhances the compound’s adsorption strength on the metal exterior side [81]. Compounds with elevated electronegativity values exhibit a pronounced ability to strongly adhere to the metal externally. The electron transfer parameter, denoted as ΔN_max, serves as a critical indicator of the interaction between a compound and a metal surface. A ΔN_max value less than zero (ΔN_max < 0) signifies electron flow from the metal surface to the molecule, while a value greater than zero (ΔN_max > 0) reflects electron transfer from the compound to the surface [82]. In the cases of both L and its complex (LFe), positive ΔN_max values confirm that the metal surface is receiving electrons from the chemicals. This electron donation plays a pivotal role in bolstering the corrosion resistance of the metal during the reaction process. Interestingly, the protective capacity of the compound increases as the ΔN_max value rises, demonstrating an enhanced ability to shield the metal surface from corrosion. FeL exhibits a ΔN_max of 4.10, surpassing the ΔN_max of L at 2.85. This result suggests that the Fe(III) complex exerts A more powerful impact in strengthening the metal’s resistance to corrosion, underscoring its superior inhibitory performance.

3.3.6. Density of States (DOS) Spectra

The GaussSum 3.0 software [83] was utilized to analyze the contributions of the groups to the molecular orbitals (HOMO and LUMO) and to generate the DOS spectrum illustrated in Figure 23. This spectrum for the L and FeL compounds was instrumental in evaluating their band gap as well as their potential anti-corrosion characteristics. The DOS spectra were obtained by blending the molecular orbital data with Gaussian curves of unit height. In the spectrum, the green and red lines represent the HOMO and LUMO levels, respectively. DOS analysis reveals that the L compound possesses HOMO and LUMO energy levels of −6.10 and −2.93 eV, respectively, corresponding to a band gap of 3.17 eV. On the other hand, the LFe complex exhibits these energy levels at −6.09 and −3.70 eV, resulting in a narrower band gap of 2.39 eV. These observations align well with HOMO-LUMO analysis, affirming the reliability of the DOS analysis for the compounds. Furthermore, the findings corroborate previous DFT results, underscoring the anti-corrosion efficacy of the FeL complex.

4. Conclusions

In acidic conditions, different concentrations of L and FeL additions are used as Ni inhibitors. It was found that the presence of inhibitors reduced the corrosion current density of nickel. The inhibition efficiency for nickel at a concentration of 1 × 10⁻³ M with L and FeL was observed to be 90.50% and 92.70% at a temperature of 25 degrees Celsius, respectively.

Several Electrochemical Impedance Spectroscopy (EIS) studies have been performed on the electrode of interest in an acidic medium, with varying concentrations of each addition. Nyquist plots showed that the presence of additives led to the development of charge transfer resistance. This behavior is observed when both L and FeL additives are used. Furthermore, Ni’s charge transfer resistance values are higher when FeL is present in different concentrations than when L is present under the same circumstances and concentrations. Additionally, when L and FeL additives are present, the double layer capacitance decreases. This result is consistent with information obtained from potentiodynamic polarization measurements of a Tafel plot.

To investigate the influence of temperature on the electrochemical and corrosive behaviors of nickel in acidic media, as well as the presence or absence of various concentrations of additives, a Tafel plot is constructed for our electrode under investigation at various temperatures. Activation energy values are calculated. The presence of a large concentration of the additive (1 × 10⁻³ M for L and FeL, respectively) could be responsible for the observed greater energy barrier, which would impede the dissolving process. The Gibb’s free energy for nickel has been calculated at various temperatures. The dissolving process becomes more spontaneous as the temperature rises, increasing the Gibbs free energy values for each addition that was present. This suggests that the inhibitor chemisorbs onto Ni’s surface. Thermodynamic parameters have been calculated. In H₂SO₄ solution, the standard enthalpy change ΔH⁰_ads Ni was calculated to be −29.14 and −31.01 kJ mol⁻¹ with varying concentrations of L and FeL. When L and FeL were present, the standard entropy change ΔS⁰_ad. values for Ni were −520 and −160 kJ/mol, respectively. The exothermic ecology of inhibitor adsorption processes was determined following the concept of ∆H^o_ads. Both negative ∆S^o_ads and negative ∆H^o_ads values are caused by the process of replacement that takes place when the additives adsorb on the surface.

Ni’s surface layers were examined using SEM analysis following Tafel plot testing. According to oxide film study data, the thin layer on the nickel surface is mostly made up of NiO and NiSO₄ in pure medium; in acidic media, the inhibiting effect of each additive concentration decreased these oxide films. In the case of utilizing the FeL additive, new particles appeared on the Ni surface, which increase with increases in the additive concentration. These species are related to FeL complex chemisorbed particles. More pure surface regions without oxides appeared in this case.

Galvanostatic charge–discharge cycling curves at applied various current densities have been constructed to study the effects of each additive on the charge–discharge features of Ni in the acidic media. It has been found that the addition of 5 × 10⁻⁵ M of FeL provides an enhanced specific capacitance of Ni at the applied current density of ±30 mA.cm⁻² in acidic media. The most enhanced discharging time values were obtained with 1 × 10⁻³ M of L present and FeL at the applied current density of 30 mA. cm⁻². The above charge–discharge results indicate that the stability of the investigated electrode increased. It was noted that the discharge time of Ni achieved its highest value in the vicinity of 1 × 10⁻³ M of L and FeL in 0.5 M H₂SO₄, specifying that its impedance behaviors are enhanced, and confirming the obtained impedance results.

Author Contributions

Conceptualization, H.M.A.-D.; methodology, H.A.E.-S.S., A.Z., R.A.E.-K., H.M.A.-D., R.M.E.-K. and A.M.A.-D.; software, M.S., R.A.E.-K., M.F.-D., M.M.A.A. and H.M.A.-D.; validation, A.Z., M.F.-D., M.M.A.A., H.M.A.-D., R.M.E.-K. and A.M.A.-D.; formal analysis, H.A.E.-S.S., M.S., R.A.E.-K., M.F.-D., M.M.A.A., R.M.E.-K. and A.M.A.-D.; investigation, H.A.E.-S.S., M.S., M.F.-D., M.M.A.A., H.M.A.-D. and A.M.A.-D.; resources, H.A.E.-S.S., M.S., A.Z., R.A.E.-K., H.M.A.-D. and A.M.A.-D.; data curation, H.A.E.-S.S., R.A.E.-K., M.M.A.A., R.M.E.-K. and A.M.A.-D.; writing—original draft, H.A.E.-S.S., M.S. and A.M.A.-D.; visualization, A.Z. and M.F.-D.; supervision, A.Z. and A.M.A.-D.; project administration, R.M.E.-K. and A.M.A.-D.; funding acquisition, H.M.A.-D., A.Z., M.S. and A.M.A.-D. All authors have read and agreed to the published version of the manuscript.

Data Availability Statement

The original contributions presented in this study are included in the article/Supplementary Materials. Further inquiries can be directed to the corresponding authors.

Acknowledgments

The authors extend their appreciation to Taif University, Saudi Arabia, through the project number TU-DSPP-2024-28. Also, the authors are thankful for the technical and human support provided by SGIker of UPV/EHU (Medidas Magnéticas Gipuzkoa) and European funding (ERDF and ESF). This research was funded by the Spanish MICIN under PID2022-141373NB-I00 project, by the EU (Horizon Europe) under “INFINITE” (HORIZON-CL5-2021-D5-01-06) and “HARMONY”, and by the Government of the Basque Country under Elkartek (MOSINCO and ATLANTIS) projects and the scheme of “Ayuda a Grupos Consolidados” (Ref.: IT1670-22).

Conflicts of Interest

The authors declare no conflict of interest.

Footnotes

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Figures, Scheme and Tables

Figure 1 ¹HNMR spectra of L imine pro-ligand.

Figure 2 Molecular electronic spectra of L ligand and its LFe complex in ethanol at 298 K.

Figure 3 Mass spectrum of (a) imine ligand L and (b) its LFe complex.

Figure 4 The stoichiometry of the produced compounds in water-based ethanolic solution using (a) Job’s and (b) molar ratio plots at 10⁻³ M of FeL complex and 298 K.

Figure 5 Tafel plot curves showing the presence of 1 × 10⁻³ M of L and Ni in 0.5 M H₂SO₄ in deficit and FeL (a), relation between corrosion rate and inhibitor concentration C_inh. in the presence of L (b) as well as FeL (c).

Figure 6 At 25 °C, the Tafel plot curves for Ni in 0.5 M H₂SO₄ with different concentrations of L (a) and FeL (b) are shown.

Figure 7 For Ni in 0.5 M H₂SO₄, with different dosages of L at 25 °C, the Nyquist plot (a), bode plot (b), and phase bode plot (c) are shown.

Figure 8 For Ni in 0.5 M H₂SO₄ with different dosages of FeL at 25 °C: Nyquist plot (a), bode plot (b), and phase bode plot (c).

Figure 9 Tafel curvatures for Nickel in 0.5 M H₂SO₄, which have 1 × 10⁻⁵ M (a), 5 × 10⁻⁵ and (b) 1 × 10⁻⁴ M (c), as well as 1 × 10⁻³ M (d) of L and 1 × 10⁻⁵ M (e), 5 × 10⁻⁵ (f) 1 × 10⁻⁴ M (g), and 1 × 10⁻³ M (h) of FeL complex at different temperatures.

Figure 10 The Arrhenius equation, when applied to the corrosion rate (I_corr) of nickel in a 0.5 M H₂SO₄ environment devoid of additives, indicates the impacts of varying temperatures for (a) L imine ligand and (b) its FeL complex.

Figure 11 The experimental findings obtained from Tafel polarization studies on Ni metal, conducted in a 0.5 M H₂SO₄ environment with varying concentrations of several inhibitors, were meticulously analyzed. The data, recorded at temperatures of 25 °C, 35 °C, 45 °C, and 55 °C, underwent fitting with the Langmuir adsorption model. This model involved the comparison of the actual surface coverage (θ) with the theoretical maximum coverage (C_inh.), considering both the presence of the inhibitors denoted by L (parameters (a–d)) and their respective FeL complexes (parameters (e–h)).

Figure 12 The graph depicts the alteration of ΔG⁰_ads as a function of temperature T, with Ni situated in 0.5 M H₂SO₄ solutions for (a) L imine ligand and (b) its FeL complex. The study incorporates multiple temperature settings to observe the effects of different inhibitor concentrations on the system.

Figure 13 SEM images of oxide films formed on the Ni surface in 0.5 M H₂SO₄, in the presence of 1 × 10⁻³ M of L (b) and in the absence of L (a), have been captured.

Figure 14 Captured SEM images for oxide films produced on Ni external side in 0.5 M H₂SO₄ in the presence of 1 × 10⁻⁵ M (a), 1 × 10⁻⁴ M (b) and 1 × 10⁻³ M (c) of FeL.

Figure 15 Analysis with EDAX for oxide layers that form on the outside of Ni when there is 1 × 10⁻⁵ M (a) and 1 × 10⁻³ M (b) of FeL present.

Figure 16 Galvanostatic charge–discharge curves of Ni electrode in 0.5 M H₂SO₄ without L (a), and with 1 × 10⁻⁵ M (b), 5 × 10⁻⁵ (c), 1 × 10⁻⁴ (d), 5 × 10⁻⁴ (e), and 1 × 10⁻³ (f) of L, at various current densities.

Figure 17 Galvanostatic charge–discharge curvatures of Ni electrode in 0.5 M H₂SO₄ containing 1 × 10⁻⁵ M (a), 5 × 10⁻⁵ (b), 1 × 10⁻⁴ (c), 5 × 10⁻⁴ (d), and 1 × 10⁻³ (e) of FeL additive at various current densities.

Figure 18 Galvanostatic charge–discharge curvatures of Ni in the absence and presence of 1 × 10⁻⁵ M (a) and 1 × 10⁻³ M (b) of L, as well as FeL additives, in 0.5 M H2SO4 at an applied current density of ±30 mA · cm⁻².

Figure 19 The suggested structure of L and FeL generated via the DFT method.

Figure 20 The MEP diagram of L and its LFe complex generated via the DFT method.

Figure 21 The Mulliken charge distribution of the L compound generated via the DFT method.

Figure 22 The distributions and HOMO and LUMO energy surface of the L and LFe complex generated via the DFT method.

Figure 23 The DFT approach produced the DOS spectrum for the L ligand and its FeL complex.

Table 1

Both physical characteristics and data from analysis of the L imine pro-ligand and its complex.

Compound (Molecular Formula)	M. Weight	Yield	Color	Molar Conductance	μeff (BM)	M.p. and Dec.temp. (°C)	Analysis (%) Found (Calcd)
				ᴧm, Ω−1 cm2 mol−1			C	H	N
L C27H18Br2N2O3	(578.25)	95%	yellow	-	-	210	56.15 (56.03)	3.18 (3.11)	4.76 (4.84)
LFe C27H26Br2FeN3O11	(784.13)	91%	bright brown	10.75	5.52	>300	41.28 (41.35)	3.25 (3.34)	5.41 (5.36)

Table 2

Ni metal Tafel constants in 0.5 M H₂SO₄ at different concentrations of the L and FeL complexes at 25 °C.

C inh.	Ecorr (V)	Icorr(µA cm−2)	βa(V/Decade)	βc(V/Decade)	Θ	η%
	In the presence of L
Blank	−0.228	616.6	19.5	−4.38	-	-
1 × 10−5 M	−0.377	371.5	23.3	−4.53	0.397	39.7
5 × 10−5 M	−0.396	223.8	26.2	−9.5	0.637	63.7
1 × 10−4 M	−0.394	125.3	31.1	−5.3	0.796	79.6
5 × 10−4 M	−0.408	64.5	33.6	−6.2	0.896	89.6
1 × 10−3 M	−0.408	58.3	33.4	−4.75	0.905	90.5
	In the presence of FeL
1 × 10−5 M	−0.387	338.3	33.4	−4.6	0.451	45.1
5 × 10−5 M	−0.384	186.4	34.5	−5.3	0.698	69.8
1 × 10−4 M	−0.370	93.0	33.1	−5.1	0.849	84.9
5 × 10−4 M	−0.358	60.2	33.5	−5.3	0.902	90.2
1 × 10−3 M	−0.344	44.67	29.3	−4.3	0.927	92.7

Table 3

Electrochemical resistance values for nickel metal in 0.5 M H₂SO₄ at 25 °C with multiple L and FeL concentrations.

In the presence of L
C inh.	Rs Ω	Rct Ω	Cdl F
Blank	1.0	13.8	1.15 × 10−6
1 × 10−5 M	1.0	17.0	9.36 × 10—7
5 × 10−5 M	1.0	19.5	8.16 × 10−7
1 × 10−4 M	1.0	24.6	6.46 × 10−7
5 × 10−4 M	5.0	39.6	4.019 × 10−7
1 × 10−3 M	5.0	43.7	3.6 × 10−7
In the presence of FeL
1 × 10−5 M	1.0	25.4	6.3 × 10−7
5 × 10−5 M	1.0	34.1	4.66 × 10−7
1 × 10−4 M	3.0	35.2	4.52 × 10−7
5 × 10−4 M	3.0	41.6	3.82 × 10−7
1 × 10−3 M	1.0	46.5	3.42 × 10−7

Table 4

Tafel parameters for the Nickel metal in 0.5 M H₂SO₄ when L is present as well as absent at various concentrations and at distinct temperatures.

With the presence of 1 × 10−5 L
Temp.	Ecorr(V)	Icorr(µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.377	371.5	23.3	−4.53	0.397	39.7
35 °C	−0.388	776.3	17.01	−4.14	0.354	35.4
45 °C	−0.391	1506.6	18.42	−4.34	0.575	57.5
55 °C	−0.389	3548.13	20.87	−4.47	0.521	52.1
With the presence of 5 × 10−5 L
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.396	223.8	26.2	−9.5	0.637	63.7
35 °C	−0.418	302	13.8	−5.6	0.75	75.0
45 °C	−0.415	432.5	28.2	−5.5	0.878	87.8
55 °C	−0.427	1738	23.3	−4.9	0.765	76.5
With the presence of 1 × 10−4 L
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.394	125.3	31.1	−5.3	0.796	79.6
35 °C	−0.393	281.8	25.9	−5.46	0.766	76.6
45 °C	−0.391	1071.6	20.81	−4.83	0.698	69.8
55 °C	−0.388	1513.6	21.1	−5.22	0.795	79.5
With the presence of 1 × 10−3 L
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.408	58.3	33.4	−4.75	0.906	90.6
35 °C	−0.392	154.8	30.1	−5.6	0.872	87.2
45 °C	−0.387	301.9	32.2	−4.7	0.914	91.4
55 °C	−0.389	703.1	32.1	−5.2	0.905	90.5

Table 5

Tafel parameters for Nickel metal in 0.5 M H₂SO₄ when FeL is present as well as absent at various concentrations and at different temperatures.

With the presence of 1 × 10−5 FeL
Temp.	Ecorr(V)	Icorr(µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.387	338.3	33.4	−4.6	0.451	45.1
35 °C	−0.385	981.7	29.5	−4.8	0.184	18.4
45 °C	−0.384	1161.4	27.6	−4.65	0.673	67.3
55 °C	−0.389	1663.4	27.9	−4.9	0.775	77.5
With the presence of 5 × 10−5 FeL
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.384	186.4	34.5	−5.3	0.698	69.8
35 °C	−0.413	407.4	28.1	−5.3	0.66	66.1
45 °C	−0.418	783.4	25.3	−5.2	0.78	78.0
55 °C	−0.414	1148.2	26.1	−6.6	0.844	84.4
With the presence of 1 × 10−4 FeL
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.370	93.0	33.1	−5.1	0.85	85.0
35 °C	−0.363	301.9	27.1	−6.3	0.748	74.8
45 °C	−0.363	297.2	28.2	−6.4	0.916	91.6
55 °C	−0.366	436.5	28.6	−6.3	0.942	94.2
With the presence of 1 × 10−3 FeL
Temp.	E corr (V)	I corr (µA cm−2)	βa(V/decade)	βc(V/decade)	Θ	η%
25 °C	−0.344	44.67	29.3	−4.3	0.927	92.7
35 °C	−0.340	89.2	28.2	−4.5	0.925	92.5
45 °C	−0.354	91.8	28.1	−4.9	0.974	97.4
55 °C	−0.348	138.1	28.8	−3.77	0.981	98.1

Table 6

The study examined the impacts of various additives on the adsorption behavior of nickel in a 0.5 M hydrochloric acid solution, analyzing the effects of different concentrations and types of additives to determine their optimal usage.

Temperature	Kads M−1 (L)	Kads M−1 (FeL)	−∆Goads. kJ·mol−1 (L)	−∆Goads. kJ·mol−1 (FeL)
298	4.4 × 104	5.9 × 103	36.44	31.4
308	2.7 × 104	6.6 × 103	36.45	32.8
318	2.2 × 104	9.02 × 103	37.11	34.6
328	1.4 × 104	1.06 × 104	37.05	36.3

Table 7

EDAX analysis of oxide films formed on the Ni metal after Tafel experiments in the presence of 1 mM of L and FeL.

L
Element	Line	Mass %	Atom %
N	K	6.22 ± 0.10	9.73 ± 0.15
O	K	32.01 ± 0.16	37.92 ± 0.22
C	K	3.16 ± 0.01	2.06 ± 0.03
S	K	8.21 ± 0.04	1.92 ± 0.03
Ni	K	50.40 ± 0.11	48.37 ± 0.09
Total		100.00	100.00
FeL
O	K	9.66 ± 0.07	15.73 ± 0.21
C	K	3.20 ± 0.01	3.05 ± 0.01
N	K	5.36 ± 0.02	7.2 ± 0.01
S	K	3.53 ± 0.03	1.23 ± 0.04
Ni	K	77.09 ± 0.11	71.73 ± 0.07
Fe	K	1.16 ± 0.01	1.06 ± 0.00
Total		100.00	100.00

Table 9

Descriptors of L and FeL compounds’ quantum reactivity.

Descriptor	L	FeL
EHOMO	−6.09	−6.08
ELUMO	−2.92	−3.69
ΔE(LUMO-HOMO)	3.17	2.39
χ	4.50	4.88
η	1.58	1.19
σ	0.63	0.84
Pi	−4.50	−4.88
ω	6.41	10.00
ΔNmax	2.85	4.10

Table 8

Charge–discharge parameters for the Ni metal in a 0.5 M H₂SO₄ solution with different concentrations of L as well as FeL at the different current densities applied.

In 0.5 M H2SO4
Blank
Applied current density	t dis.	ΔV	S.C. (mA. hr−1)
1	9.15	0.037	0.068
3	9.7	0.104	0.077
5	10.1	0.154	0.091
10	11.1	0.261	0.118
15	11.0	0.344	0.133
20	11.6	0.412	0.156
30	11.6	0.531	0.182
In the presnce of 1 × 10−5 M of L
Applied current density	t dis.	ΔV	S.C.
1	8.5	0.034	0.069
3	8.8	0.107	0.685
5	9.2	0.184	0.069
10	9.6	0.313	0.085
15	10.4	0.414	0.105
20	11.3	0.51	0.123
30	11.8	0.668	0.147
In the presence of 5 × 10−5 M of L
Applied current density	t dis.	ΔV	S.C.
1	9.0	0.035	0.714
3	9.2	0.119	0.064
5	10.2	0.177	0.08
10	10.6	0.28	0.105
15	11.2	0.353	0.132
20	11.3	0.148	0.422
30	11.4	0.174	0.544
In the presence of 1 × 10−4 M of L
Applied current density	t dis.	ΔV	S.C.
1	10.48	0.046	0.063
3	10.8	0.129	0.069
5	11.1	0.281	0.054
10	11.6	0.349	0.092
15	12.0	0.349	0.14
20	12.8	0.42	0.17
30	12.8	0.543	0.197
In the presence of 5 × 10−4 M of L
Applied current density	t dis.	ΔV	S.C.
1	9.8	0.079	0.0344
3	10.7	0.197	0.045
5	11.75	0.265	0.062
10	12.4	0.373	0.092
15	12.7	0.467	0.113
20	13.1	0.569	0.128
30	13.3	0.736	0.151
In the presence of 1 × 10−3 M of L
Applied current density	t dis.	ΔV	S.C.
1	9.8	0.096	0.028
3	11.5	0.22	0.043
5	11.7	0.28	0.568
10	12.0	0.42	0.079
15	12.6	0.532	0.098
20	13.0	0.639	0.113
30	13.2	0.822	0.134
In the presence of FeL
In the existence of 1 × 10−5 M of FeL
Applied current density	t dis.	ΔV	S.C.
1	10.6	0.074	0.039
3	11.0	0.194	0.047
5	12.0	0.266	0.063
10	12.3	0.383	0.089
15	12.8	0.475	0.112
20	12.9	0.549	0.131
30	12.9	0.690	0.155
In the presence of 5 × 10−5 M of FeL
Applied current density	t dis.	ΔV	S.C.
1	9.0	0.055	0.045
3	10.0	0.164	0.051
5	10.7	0.228	0.065
10	10.9	0.342	0.088
15	12.0	0.403	0.124
20	12.5	0.481	0.144
30	12.8	0.593	0.179
In the presence of 1 × 10−4 M of FeL
Applied current density	t dis.	ΔV	S.C.
1	9.8	0.079	0.034
3	11.0	0.187	0.049
5	11.1	0.265	0.058
10	12.1	0.388	0.086
15	12.6	0.433	0.121
20	12.8	0.644	0.110
30	13.1	0.644	0.170
In the presence of 5 × 10−4 M of FeL
Applied current density	t dis.	ΔV	S.C.
1	8.7	0.078	0.031
3	9.9	0.213	0.038
5	10.5	0.29	0.05
10	11.1	0.419	0.073
15	11.4	0.521	0.091
20	11.9	0.651	0.102
30	12.8	0.807	0.132
In the presence of 1 × 10−3 M of FeL
Applied current density	t dis.	ΔV	S.C.
1	9.82	0.073	0.037
3	10.5	0.206	0.042
5	11.1	0.299	0.052
10	11.7	0.465	0.698
15	12.0	0.596	0.084
20	12.6	0.303	0.231
30	13.3	0.690	0.161

Supplementary Materials

The following supporting information can be downloaded at: https://www.mdpi.com/article/10.3390/met15040403/s1, Figure S1: ¹³C NMR spectra of the L imine pro-ligand; Figure S2: proposal structure of L imine ligand; Figure S3: FT-IR spectrum of L imine ligand and its complex Figure S4: EDAX analysis for oxide layers formed on Ni exterior side in 0.5 M H₂SO₄; Figure S5: Change of capacity during six cycles for Ni metal in 0.5 M H₂SO₄ (a) without and with 1 × 10⁻⁵ M (b) 5 × 10⁻⁵ M (c), 1 × 10⁻⁴ M (d), 5 × 10⁻⁴ M (e), and 1 × 10⁻³ M (f) of L at applied various current densities; Figure S6: Power variation of Ni metal during six cycles in 0.5 M H₂SO₄ (a) without and with 1 × 10⁻⁵ M (b) 5 × 10⁻⁵ M (c), 1 × 10⁻⁴ M (d), 5 × 10⁻⁴ M (e), and 1 × 10⁻³ M (f) of L at applied various current densities; Figure S7: Change of Energy with cycle number for Ni metal in 0.5 M H₂SO₄ (a) without and with 1 × 10⁻⁵ M (b) 5 × 10⁻⁵ M (c), 1 × 10⁻⁴ M (d), 5 × 10⁻⁴ M (e), and 1 × 10⁻³ M (f) of L at applied various current densities; Figure S8: Change of Capacity during six cycles for Ni metal in 0.5 M H₂SO₄ with 1 × 10⁻⁵ M (a) 5 × 10⁻⁵ M (b), 1 × 10⁻⁴ M (c), 5 × 10⁻⁴ M (d), and 1 × 10⁻³ M (e) of FeL at applied various current densities; Figure S9: Power variation during six cycles for Ni metal in 0.5 M H₂SO₄ with 1 × 10⁻⁵ M (a) 5 × 10⁻⁵ M (b), 1 × 10⁻⁴ M (c), 5 × 10⁻⁴ M (d), and 1 × 10⁻³ M (e) of FeL at applied various current densities; Figure S10: Change of Energy with time for Ni metal in 0.5 M H₂SO₄ with 1 × 10⁻⁵ M (a) 5 × 10⁻⁵ M (b), 1 × 10⁻⁴ M (c), 5 × 10⁻⁴ M (d), and 1 × 10⁻³ M (e) of FeL at applied various current densities; Scheme S1: Fragmentation pathway of LFe complex. Below each structure, the precise masses of the particles and their chemical formulas are displayed; Table S1: Tafel parameters for Nickel metal in 0.5 M H₂SO₄ without additives at distinct temperatures; Table S2: Activation energy data (E_a) of Ni in [0.5 M H₂SO₄] solution, while additives are absent and present in multiple doses; Table S3: EDAX analysis for the oxide layers formed on Ni surface in 0.5 M H₂SO₄.

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