Reactivity trends of hydroxide ion attack on high spin Fe(II) complexes including bromosalicylidene amino acid ligands in some mixed aqueous solvents: Gibb’s Free Energy of Transfer and initial-transition state analysis

1.1 Application of Schiff base amino acid complexes

From a bioinorganic point of view, iron Schiff base complexes provide useful structural and electronic models for the similarly coordinated sites found in the heme iron enzymes. Moreover, these complexes are also important for the asymmetric oxidation of organic substrates, since their structure and catalytic activity are analogous with those of iron porphyrins (Canali and Sherrington, 1999). Schiff base amino acid complexes act as good chelating agents (Nath and Yadov, 1997; Abdel-Rahman et al., 2014) and behave as efficient biologically active (Abdel-Rahman et al., 2013; El-Said et al., 2001) and cytotoxic (Wang et al., 2002) agents. In addition, Schiff base amino acid complexes are considered to constitute new kinds of potential antibacterial and anticancer reagents (Wang et al., 2005).

1.2 Medium effects on reactivity

The effects of both solvents upon a variety of reactions of transition metal complexes have been of interest in recent years. Most reactions of inorganic complexes exhibit different reactivities in micellar and microemulsion systems from those in aqueous solutions. Thus, the appropriate use of microemulsions can allow kinetic studies to be made of reactions involving either water-insoluble reactants or, in some cases, such Schiff base amino acid complexes, hydrophobic substrates (Burgess and Pelizzetti, 1988; Burgess and Pelizzetti, 1992). More recently, trends in transfer chemical potentials and their relation to reactivity trends for substitution and redox reactions of these inorganic complexes have been investigated in binary aqueous mixtures (Blandamer et al., 1988). Furthermore, a new approach to understanding solvent effects on reaction rates involves consideration of the extent of the interaction of the solvent environments of the initial state and the transition state with the solvent co sphere of an added substance (Milde et al., 1999). Most treatments of solvent effects on reactivity deal with the rate constant and activation parameter trends, but others have recognized that such kinetic parameters are composite values. Analysis of solvent effects on reactivities in terms of the contributions from the initial and transition states can lead to full explanations of these trends (Burgess and Pelizzetti, 1988; Burgess and Pelizzetti, 1992; Abu-Gharib et al., 2011). In parallel with these trends, solvation effects can be probed via the variation of activation volume with the properties of the solvent (Bähr and Döge, 1957). These two approaches are complementary in that analysis of the initial and transition states provides information on solvation changes for a reaction involved in a suggested mechanism on transfer from one medium to another, whereas knowledge of the activation volume can test solvation changes on going from the initial to the transition state in a specific medium (Yamagishi, 1986; Abu-Gharib et al., 2011; Abu-Gharib et al., 2017). It is worthy to recall that thermodynamic functions, e.g., solubility, provide information on transfer of the initial state, whereas solvent effects on the transition state can be derived from these measurements and the observed rate constants (Bähr and Döge, 1957). In the present contribution, we present an extensive investigation of these approaches through the study of solvent effects on the reactivity of the base hydrolysis of a new series of hydrophobic high-spin iron(II) amino acid Schiff base chelates. Moreover, the effect of solvent on reactivity trends has been analyzed into initial and transition state analysis to give an overall idea about the role of solvent in the reactivity.

2.1 Preparation of the investigated complexes

The solid Schiff base amino acid Fe(II) complexes were prepared by mixing aqueous solutions of the respective amino acids [l-Alanine (ala), l-Phenylalanine (phala), l-Aspartic acid (asp), l-Histidine (his) and l-Arginine (arg)] with an equimolar hot ethanolic solution of 5-bromosalicyaldhyde (BrSal). The resulting Schiff base amino acid ligands (the ligands are displayed and labeled in Scheme 1) were then treated with an aqueous solution of ferrous ammonium sulfate in an equimolar ratio. In order to avoid the oxidation of Fe(II) and the formation of Fe(OH)₃, a few drops of glacial acetic acid were added (Sharma and Dubey, 1994; Abdel-Rahman et al., 2013, 2014; Blandamer et al., 1992). The resulting solution was stirred for 8 h under nitrogen. The precipitated complexes were filtered off, washed with a water–ethanol mixture, and then with diethyl ether. The complexes were first dried on a water bath and then in vacuo over P₂O₅. The isolated complexes were crystallized from water–ethanol solutions (Abdel-Rahman et al., 2014). The composition of the complexes was established by CHN microanalysis, IR, and UV/vis spectral analyses. All gave satisfactory CHN microanalysis (Abdel-Rahman et al., 2014). The stability of solutions was tested for at least a month for evidence of oxidation of the Fe(II) cation by peroxopersulfate or reduction by dithionite. When the aged complex solutions were treated with NaOH under nitrogen, a green precipitate of Fe(OH)₂ formed, indicating the maintenance of Fe(II) oxidation state in the complex solutions used in the kinetic runs. It was observed that, in these runs, the intensely violet color of the complex fades and the solution turns colorless. Some green colloidal particles of Fe(OH)₂ form initially, turning pale yellow, and, finally, precipitating as brown Fe(OH)₃ by oxidation with O₂ dissolved in solution kept long after end of the kinetic run. Full details of the characterization of the present complexes can be found in our previous publication (Abdel-Rahman et al., 2014).

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Scheme 1 Structures and abbreviations of the Schiff base ligands and abbreviations of their corresponding complexes.

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2.2 Kinetics

Kinetic measurements were carried out spectrophotometrically by following the decrease in absorbance with time at λ_max in the UV–vis region appropriate to each complex (see Scheme 2). The kinetic data were recorded using PG UV–Visible spectrophotometer model T + 80 with 10 mm matched quartz cells connected with an ultrathermostate (CRIOTERM model 190) water circulator. The reactants were brought together so that the reaction exhibits pseudo-first order kinetics by mixing a multifold greater concentration of the base than that of the complex. It was confirmed that there is no interference from other reagents at the selected wavelength absorption maxima for the investigated complexes. Rate constants were calculated from the dependence of absorbance on time at λ_max for the base hydrolysis of each complex. Each experiment was duplicated to ensure satisfactory reproducibility of the kinetic parameters (Bähr and Thämlitz, 1955) (see Scheme 3).

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Scheme 2 Scheme of pathway for the base hydrolysis of the investigated complexes.

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Scheme 3 The Suggested mechanism for the base hydrolysis reaction of the investigated complexes.

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2.3 Solubilities

Solubilities were measured at 298 K ± 0.1 K by agitating a generous excess of each complex with the appropriate water–ethanol, water–acetone, water–propanol and water–dimethyl sulfoxide mixtures in a thermostated vessel for 7 h. This time is enough to prepare the best saturated solution from the solid compound in aqueous-solvent mixtures. Then portions of supernatant saturated solution were removed and centrifuged rapidly using centrifuge before withdrawing an aliquot by using a micropipette. The solution was diluted as necessary. Then the absorbance was measured at λ_max for each complex (Bähr and Thämlitz, 1955; Blandamer et al., 1986). All aliquots of equilibrated solutions were diluted with the solvent and the dependences of λ_max on solvent composition, were found to be very small for the studied complexes, and thus ignored in calculation of solubilities. The customary precautions were taken to avoid heating of aliquots of saturated solutions before dilution and to prevent any solid material being carried through with samples of saturated solutions. The spectra were measured at different sampling times to verify that a constant concentration has been achieved. The concentrations of the saturated solutions were determined by applying beer’s law (A_abs = εcl) where “ε” is the molar extension coefficient in dm³ mol⁻¹ cm⁻¹, “c” is the concentration in mol dm⁻³ and “l” is the length of the cell in cm. The molar extension coefficient “ε” in dm³ mole⁻¹ cm⁻¹ for each complex was determined from the least square of the standard curve using beer’s law.

3.1 Kinetic results and reaction mechanism

The fading of color of the solutions of the Fe(II) amino acid Schiff base complexes by hydroxide ion indicates that the complete dissociation occurs. In the presence of excess hydroxide ion and oxygen, the ultimate product, in addition to the ligand, is Fe(OH)_3. The overall rate law for the investigated reaction under the adopted conditions of pseudo first order kinetics can be represented as follows:

The values of observed rate constants for the base hydrolysis reaction of bsphali were compared with the corresponding derivative nphali from previous publication (Abdel-Rahman et al., 2013) (cf. Table 2). The data show that the effect of function group on the reactivity trends of the prepared complexes toward the attack of OH⁻ ion on them. It is observed that the reactivity of bsphali toward the base hydrolysis is more than the corresponding nphali. This is because the inductive effect of bromide atom in bsphali is more than the phenyl ring in nphali.

From the comparison between the values of second-order rate constant (k₂) of the base hydrolysis of the studied complexes (cf. Table 3), it was found that the values of k₂ are correlated with the effect of side chain substituent R in the structure of the complexes under investigation. The order of reactivity of the prepared bromosalicylidene complexes toward the hydroxide attack is increased in the sequence: bsari < bsali < bsphali < bshi < bsasi. This may be rationalized to the inductive effect of the substituent. This may be due to the electron withdrawing properties of carboxylic group more in bsasi than imidazole ring in bshi and phenyl ring in bsphali. Where in bsari, the presence of a guanidine group which is strong electron donating group makes its reactivity toward OH⁻ low.

The solubility of the prepared complexes was measured in aqueous solution, aqueous-organic solvents such as ethanol, propanol, acetone and dimethyl sulfoxide (cf. Table 4). The transfer chemical potentials for an uncharged compound are generally derived from solubility measurement on the assumption that the ratio of mean activity values co-efficient in the aqueous and aqueous-organic mixtures was in all cases unity (Bähr and Thämlitz, 1955) according to the following thermodynamic equation:

The hydrophobicity of the investigated Fe(II) Schiff base amino acid complexes can be rationalized to the low solubility of the investigated complexes in aqueous medium (cf. Table 4). Moreover, the results of negative values transfer chemical potentials of none charged the investigated Fe(II) Schiff base amino acid complexes (cf. Tables 5–9) can mainly be ascribed to the hydrophobic nature of the investigated complexes. It is observed that the hydrophobicity of the investigated complexes leads to an increase in the stabilization of them with increasing solvent percent. The order of the hydrophobicity increases in the following order: $$\text{bsasi}<\text{bsali}<\text{bsari}<\text{bshi}<\text{bsphali}$$ Thus the results show preferential solvation by the solvent as molecule size increases.

The effect of co-organic solvent on the reactivity has been investigated in the presence of ethanol, propanol, acetone and dimethyl sulfoxide and the observed first-order rate constant values (k_obs) for the base hydrolysis were evaluated and reported in Table 4 and Figs. 1 and 2. The increase of the added co-solvent % exhibits a general decrease of the rate constants of the base hydrolysis of the examined complexes. This reactivity trend is attributed to the opposed effects of the extensive destabilization or to the salting out of the OH⁻ ion (Abu-Gharib et al., 2011; Abu-Gharib et al., 2017; El-Khatib and Nassr, 2007). Furthermore, the high delocalization of charge within Fe(II) azomethine complex promotes interaction with localized dispersion centers in nearby solvent molecules (Abu-Gharib et al., 2011, 2017; Shaker and Nassr, 2002). Thus, the destabilization of OH⁻ ions and conversely the stabilization of the complex in the presence of increasing amounts of co-solvent in the reaction aqueous medium, as evidenced from the remarkably increased solubility on going from the aqueous to 60% (v/v) co-solvent solution as shown in Tables 4 and 5 reduce the opportunity of the reactants to combine together to afford the activated complex and thus reduces the reaction rate. The effective density of dispersion centers in solvent molecules increases in the order H₂O < DMSO < acetone < propanol < ethanol. Thus this type of interaction is expected to be weak as the mole fraction of water increases. Thus, in turn, leads to destabilize the azomethine compound in the same direction.

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Figure 1 Plots of the observed first-order rate constant values for the base hydrolysis of bsphali complex in different ratios (v/v) of solvent at [bsphali] = 1.0 × 10⁻⁴ mol dm⁻³, [OH⁻] = 3.33 × 10⁻³ and 298 K.

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Figure 2 Plots of the change in the activation barrier (δ_mΔG^#) for the base hydrolysis of bsphali complex in different ratios (v/v) of solvent at 298 K.

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In accordance with the proposed mechanism in all of the systems, the transition state is more hydrophilic than the initial state, because of the more ionic character of the former. In general, the transition state would be less stabilized by solvation with increasing amounts of the co-solvent compared with the initial state. This effect, combined with the destabilization of the convenient hydrophilic simple OH⁻ ion, would lead to an enhanced reaction rate due to the greater opportunity of the destabilized reactants to bind. However, this is not the case for the recently observed reactivity trends in most complexes investigated. This paradox in the reactivity trends would be intuitively attributable to competition between the opposing effects of the stabilization by dispersion or solvation of the particular complex, on one hand, and the destabilization of OH⁻ ions in co-solvents, on the other hand. If compensation between these two effects occurred, there should be small effects on reactivity on adding organic co-solvents to the aqueous reaction.

Tables 5–9 and Figs. 2 and 3 indicate that the values of activation energy barrier (δ_mΔG^#) at different ratios of co-solvent increase with increase in ethanol, propanol, acetone and dimethyl sulfoxide %. These trends match with the decrease in the values of k₂ as co-solvent % increases.

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Figure 3 Plots of the change in the activation barrier (δ_mΔG^#) for the base hydrolysis of bsasi complex in different ratios (v/v) of solvent at 298 K.

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3.2 Initial state-transition state analysis for the investigated complexes in water-solvent mixtures

Solvent effects on the rate constants reflect solvation changes on the initial and transition states as these are transferred from water into mixed aqueous methanol or aqueous acetone media. If solvent effects on the initial state can be established from thermodynamic data (solubility data), then solvent effects on the transition state can be calculated from initial state and kinetic data assuming constancy of mechanism and validity of transition state theory. The free energies of activation and transfer chemical potentials are now used in our analysis of reactivity trends with changing solvent composition into initial state and transition state components, i.e., combination of kinetic data, solubility data and transfer chemical potentials yield the effect of solvent on the transition state according to equation (Abu-Gharib et al., 2011 ,2017; Blandamer et al., 1986; El-Khatib, 2002):

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Figure 4 Plots of the initial state (IS) and transition state (TS) for the base hydrolysis of bsphali complex in different ratios (v/v) of EtOH at 298 K.

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Figure 5 Plots of the initial state (IS) and transition state (TS) for the base hydrolysis of bsphali complex in different ratios (v/v) of Ac at 298 K.

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Figure 6 Plots of the initial state (IS) and transition state (TS) for the base hydrolysis of bsphali complex in different ratios (v/v) of DMSO at 298 K.

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