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Original article
12 (
8
); 4240-4250
doi:
10.1016/j.arabjc.2016.05.007

Mechanism of the oxidation of 1-(ferrocenyl)-ethanone/ethanol by dicyanobis(phenanthroline)iron(III)

, , ,
a Department of Chemistry, Shaheed Benazir Bhutto Women University, Peshawar, Pakistan
b Department of Chemistry, University of Karachi, Karachi 75270, Pakistan
c Department of Chemistry, Jinnah University for Women, V-C Nazimabad, Karachi, Pakistan
d Department of Chemistry, COMSATS Institute of Information Technology, 22060 Abbottabad, Pakistan

⁎Corresponding author at: Department of Chemistry, Shaheed Benazir Bhutto Women University, Peshawar, Pakistan. rznkhattak@sbbwu.edu.pk (Rozina Khattak) rznkhattak@yahoo.com (Rozina Khattak)

Received: , Accepted: ,
© The Author(s) 2016
Disclaimer:
This article was originally published by Elsevier and was migrated to Scientific Scholar after the change of Publisher.

Peer review under responsibility of King Saud University.

Abstract

This study highlights that dicyanobis(phenanthroline)iron(III); [FeIII(phen)2(CN)2]+ oxidises 1-ferrocenylethanone; [CpFeIICpCOMe], and 1-ferrocenylethanol; [CpFeIICpCHOHMe] through a complex mechanism. The kinetic data suggest that the reaction(s) whilst undergoing completion passed through three stages. The first stage was noted soon after starting the reaction(s) and that the reaction(s) followed an overall zeroth order in the first stage that continued to a ∼30% and ∼51% proceeding of the reaction in case of 1-ferrocenylethanone and 1-ferrocenylethanol, respectively. The reaction(s) was found to follow overall second order in the second phase. The second phase culminated into a third phase near the end of the reaction(s) when the rate of redox reaction(s) competed the rate of decrease in solubility of one of the product species; [FeII(phen)2(CN)2]. The third phase may be termed as the competition phase. We identified these phases because the reactants and the intermediates participated in the reaction(s). Protonation of the reducing agents generated intermediates. The intermediates reduced [FeIII(phen)2(CN)2]+ to [FeII(phen)2(CN)2] with a different order of reaction as compared to the initial reactants. The effect of variation in the concentration of protons (0.2–5.2 mM) showed that the conjugate acid of 1-ferrocenylethanone; [CpFeIICpC+OHMe] catalyses the oxidation of [CpFeIICpCOMe]. The protonation of 1-ferrocenylethanol to form [CpFeIICpCHO+H2Me] reduced the rate of oxidation of [CpFeIICpCHOHMe]. These opposing results were verified by observing that the rate of oxidation of [CpFeIICpCOMe] was increased by an elevation in the ionic strength (primary salt effect), and deaccelerated by reducing the dielectric constant or increasing the proportion of dioxane, 10–50 (% v/v). Together, these results support the pivotal role of [CpFeIICpC+OHMe] in the rate-determining step. Neutral [CpFeIICpCHOHMe] takes part in the rate-determining step of the redox reaction between [CpFeIICpCHOHMe] and [FeIII(phen)2(CN)2]+, where the studies pertaining to influence of ionic strength and dielectric constant corroborate these findings. These outcomes of the study helped to postulate a mechanism for each reaction, and to propose the rate law(s).

Keywords

Iron coordination complexes
Dicyanobis(phenanthroline)iron(III)
1-Ferrocenylethanone
1-Ferrocenylethanol
1,4-Dioxane
Visible absorption spectrophotometry
1

1 Introduction

The organometallic compounds, 1-ferrocenylethanone and 1-ferrocenylethanol, although well known, have been little studied for their electron transfer tendencies with other transition metal complexes. The potential of such studies is underlined by the antineoplastic activity of ferrocene-containing alcohols (Shago et al., 2007). The activity of an antineoplastic drug relies on the ferrocenium species, which is FeIII part of the drug. Redox-active enzymes oxidise the ferrocenyl group (FeII form) of the drug to the ferrocenium (FeIII) species, which aids in Chemotherapy (Osella et al., 2000). The mechanism generates hydroxyl radicals (OH•) through the interaction of the resulting ferrocenium species with water and oxygen. DNA strands are cleaved by hydroxyl radicals resulting in cell death. Studies of the one electron-exchange electrochemistry of the derivatives of ferrocene reveal important facets of these compounds in medication as well as in other procedures. We studied the oxidation of 1-(ferrocenyl)-ethanone/ethanol by dicyanobis(phenanthroline)iron(III) in aqueous dioxane. The study aims to follow the reactivity profiles of the two complexes and to compare them, and also to explore mechanistic pathways.

A literature survey of the reactions of 1-ferrocenylethanone (E1/2 0.74 V) and 1-ferrocenylethanol (E1/2 0.204 V) identified various derivatives and their electrochemical characteristics, but thorough kinetic studies that highlight the mechanism(s) were absent (Baciocchi et al., 1993; Bechki and Lanez, 2010; Casas et al., 2007; Glidewell et al., 1996; Hamilton, 1991; Jong et al., 1999; Li et al., 2011; Ryabov et al., 1999; Sasaki and Pittman, 1973; Wang and Anzai, 2015; Xu et al., 1994). The oxidation of 1-ferrocenylethanone and 1,1′-diacetylferrocene has been studied by ceric sulphate in sulphuric acid (Holeček et al., 1979). This study suggests mechanisms, but neglects the probable formation of various ceric complexes, which may take part in the reactions, as we described in our earlier work (Khattak et al., 2008). Molecular oxygen has also been used to oxidise 1-ferrocenylethanone and formylferrocene in organic solvents, recently (Fomin and Shirokov, 2012). The comprehensive kinetic measurements helped to propose different mechanisms. We extended this approach and use a mixed-ligand FeIII complex acting as a strong oxidising agent (E1/2 0.80 V). This system has been used against a number of compounds of biological importance (Blake et al., 1991; Matsumoto et al., 1997; Pelizzetti et al., 1978; Takagi et al., 1995). Our interest in this oxidant revolves around its few distinctive properties such as its solubility in mixed aqueous–organic solvent mixtures, high reduction potential and highly coloured reduced (FeII) form. These characteristics also helped us to follow the reaction kinetics spectrophotometrically.

2

2 Experimental

Analar grade materials; potassium nitrate, nitric acid, 1,4-dioxane, ferrous ammonium sulphate hexahydrate, 1,10-phenanthroline, were used without further purification, unless otherwise stated. 1-Ferrocenylethanone and 1-ferrocenylethanol (95–97%) were purchased from Acrōs Organics. A nitrate salt of dicyanobis(phenanthroline)iron(III) was synthesised, and purified subsequently, according to reported procedures (Khattak, 2011; Schilt, 1960). The methods and results are briefly explained in the details that follow. We used distilled and deionised water in our kinetic experiments.

2.1

2.1 Synthesis of [FeIII(phen)2(CN)2]NO3⋅2H2O

The synthesis of FeII complex, [FeII(phen)2(CN)2]⋅2H2O is required to synthesise the mixed-ligand FeIII complex, [FeIII(phen)2(CN)2]NO3⋅2H2O. The multistep synthesis involved the preparation of homoligand chelate, tris(phenanthroline)iron(II). This yielded [FeII(phen)2(CN)2], upon further treatment with cyanide. The dihydrate compound was crystallised out and dried, and then after oxidised with concentrated nitric acid and purification, the nitrate salt of the FeIII complex was obtained.

Ferrous ammonium sulphate hexahydrate was dissolved with 1,10-phenanthroline in warm water in a mole ratio; 1:3. The solution was heated and stirred continuously to ensure complete dissolution and the reaction of phenanthroline with FeII salt to form the complex, [FeII(phen)3]2+. A fresh concentrated aqueous solution of potassium cyanide (small volume) was added, all at once, to the near boiling solution with constant stirring. Small dark violet crystals of dicyanobis(phenanthroline)iron(II) settled down slowly, after moving the hot vessel to an ice-water bath. The vessel was moved to a dry place and stored at ambient temperature after ∼55 min. Three days were long enough for ageing. The crystals were collected by suction filtration, washed with water, and dried. This step can be represented by the following equation:

(1)

The crystalline FeII complex (1 g) was oxidised to the FeIII complex by slow heating and continuous stirring with concentrated nitric acid (4.5 mL). Reddish-brown nitrogen dioxide (NO2) gas was evolved during this process. Evolution of the gas stopped on completion of the reaction. The hot purple solution was then diluted to 36 mL with water. Finely divided purple precipitates appeared, which were re-dissolved by heating to yield a clear purple solution. On cooling at room temperature and overnight standing dark purple glistening needles slowly developed. After a few days ageing, filtration, washing with water, and drying yielded crude FeIII complex, which required passing through the purification process detailed ahead. Following equation summarises the whole process:

(2)

The purification process involved further treatment of the resulting crystalline product with 6–7 mL of concentrated nitric acid followed by warming and stirring continuously. Upon dilution to 30 mL with water and 46–47 mL with 1,4-dioxane, the solution was heated for a further 10 min. After cooling to room temperature product recrystallised over ∼45 h. Filtration and washing with water, drying, and elemental analysis confirmed the dark purple glistening needles to be [FeIII(phen)2(CN)2]NO3⋅2H2O. Elemental analysis of [Fe(phen)2(CN)2]NO3·2H2O (FW 566.27): %Calculated: Fe, 9.86; C, 55.10; H, 3.53; N, 17.30. %Found: Fe, 9.85; C, 55.79; H, 3.20; N, 17.27. Average percent purity: 97.91.

2.2

2.2 Kinetic measurements

Kinetic studies were performed under pseudo-first order condition, and at a constant pH, ionic strength, dielectric constant, and temperature. The concentration of reducing agent was always in excess over the oxidising agent. In order to identify the influence of each of these individual factors upon the rate of reaction, the reaction was probed by retaining all other components constant. Formation of the reduced dicyanobis(phenanthroline)iron(II); [FeII(phen)2(CN)2], was followed spectrophotometrically at 530 nm in 80% (v/v) aqueous dioxane (supplementary information, Figs. S1 and S2), where dioxane co-solvent was required because of the sparingly soluble nature of 1-(ferrocenyl)-ethanone/ethanol in water. The values of molar absorptivity (ε) of [FeII(phen)2(CN)2] and [FeIII(phen)2(CN)2]+ were determined and found to be 10,273 M−1 cm−1 and 1007 M−1 cm−1, respectively in 80% aqueous dioxane at 530 nm. The stock solution(s) of 1-(ferrocenyl)-ethanone/ethanol was prepared in 1,4-dioxane, which was diluted further with water or the solvent mixture to obtain the required percentage. Low concentrations of the complex were able to produce good absorbance responses according to Beer–Lambert Law, as the changes in absorbance with respect to time were monitored experimentally. We did not observe any considerable interference in absorbance of [FeII(phen)2(CN)2] by the reactants ([FeIII(phen)2(CN)2]+, [CpFeIICpCOMe] and [CpFeIICpCHOHMe]) and the products ([CpFeIIICpCOMe]+ and [CpFeIIICpCHOHMe]+) during the kinetic studies. The absorbance at the zero time point was subtracted from the absorbance at a time point ‘t’, at the higher concentrations of 1-(ferrocenyl)-ethanone/ethanol despite blank correction using solutions that were devoid of [FeIII(phen)2(CN)2]+ system. This approach produced good data with respect to [FeII(phen)2(CN)2]. Each experiment was repeated ca. 6 times in order to check accuracy. Fresh solutions of the oxidant and other reactants were used in each experiment. Significant care was always taken to store solutions in a dry and dark place at ambient temperature before making use of them, bearing in mind the report of Papula et al. (1990) on the photoinduced reduction of [FeIII(bpy)2(CN)2]+. We found no evidence of instability of the oxidant and other reactants during the course of these kinetic studies in lieu with the findings of Blake et al. (1991) and Wang and Stanbury (2008). We did not observe increase in the absorbance after starting the “experiment” and before initiating the “reaction” by adding reductant into the cuvette (supplementary information, Fig. S3 does not show any photoinduced reduction of [FeIII(phen)2(CN)2]+).

2.3

2.3 Instrumentation

The selected electron transfer processes followed fast kinetics (in order of milliseconds). A ‘home-built’ rapid reaction monitoring system was employed for data acquisition. The detailed description of the instrumental assembly is given in the supplementary information.

The spectral characterisation of the compounds was performed on a Shimadzu UV-160A UV/Vis spectrophotometer, and elemental analysis on Carlo Erba 1106 elemental analyser. IKA Combimag RCH accessory was used for routine heating and stirring. Mettler Toledo MP-220 basic pH/mV/°C meter, and HANNA HI 8314 membrane pH meter equipped with SCHOTT pH-electrode Blueline 25PH (pH 0–14/−5 to +80 °C/Gel) were used for pH measurements. Haake KT 33 cooling water bath with heater and circulator helped to maintain temperature.

3

3 Results and discussion

The kinetic measurements demonstrate that each reaction was completed in three phases (Figs. 1 and 2). In the first phase of the reaction(s), as the reaction(s) starts, an overall zeroth order was observed: a zero-order reaction with respect to [CpFeIICpCOMe]/[CpFeIICpCHO+H2Me], and [FeIII(phen)2(CN)2]+ each. Almost 30% and ∼51% fraction of the reaction of 1-ferrocenylethanone and 1-ferrocenylethanol with dicyanobis(phenanthroline)iron(III) respectively underwent a zeroth order. In the second phase, the reaction(s) is first order with respect to each [CpFeIICpC+OHMe]/[CpFeIICpCHOHMe], and [FeIII(phen)2(CN)2]+. An overall second order reaction was thus observed. The second phase reached to its end as the reaction completed, but due to the saturation of the reaction mixture with the neutral product, [FeII(phen)2(CN)2] near the end of the reaction, we observed a third phase. We monitored increase in absorbance as a function of time during data acquisition. The third phase was recognised near the end of the reaction, when the concentration of product; [FeII(phen)2(CN)2] was maximum. The presence of precipitates of [FeII(phen)2(CN)2] in the cuvette after completion of the reaction indicated decrease in its solubility in aqueous dioxane. Having presumed, the low solubility of [FeII(phen)2(CN)2] in aqueous dioxane at its higher concentrations due to its electrical neutrality, the solubility drop proved to become a factor that altered the rate of redox reaction(s). The partition in phases constrained to use integrated rate equations to determine the rate constants. Absorbance indeed depends upon the concentration of [FeII(phen)2(CN)2], according to Beer’s law, the plot(s) between ln [AY]t on y- and time on x-coordinate was drawn to the maximum appropriate time points before linearity of the data bent down in order to find out the observed pseudo-first order rate constant(s); (k′obs), and to circumvent the problem. However after bending, the plot(s) still retained linearity, but yielded a comparatively different value(s) of the slope(s), pseudo-first order rate constant, which highlighted the third phase as a competition phase. The plots were always linear with R2 1.0 (supplementary information, Figs. S5–S8). The observed rate constants, which were acquired from the slopes of the plots, are the average of 3–6 values.

Kinetic trace of the reaction between 1-(ferrocenyl)-ethanone/ethanol and dicyanobis(phenanthroline)iron(III) in 80% (v/v) aqueous dioxane at 530 nm, and 0.18 mM ‘I’. Plot “A” represents different phases of the reaction between 1.3 mM [CpFeIICpCOMe] and 0.08 mM [FeIII(phen)2(CN)2+] at 304 ± 0.5 K. Plot “B” is the reaction profile of 1.3 mM [CpFeIICpCHOHMe] with 0.08 mM [FeIII(phen)2(CN)2+] at 291 ± 0.5 K.
Figure 1 Kinetic trace of the reaction between 1-(ferrocenyl)-ethanone/ethanol and dicyanobis(phenanthroline)iron(III) in 80% (v/v) aqueous dioxane at 530 nm, and 0.18 mM ‘I’. Plot “A” represents different phases of the reaction between 1.3 mM [CpFeIICpCOMe] and 0.08 mM [FeIII(phen)2(CN)2+] at 304 ± 0.5 K. Plot “B” is the reaction profile of 1.3 mM [CpFeIICpCHOHMe] with 0.08 mM [FeIII(phen)2(CN)2+] at 291 ± 0.5 K.
Kinetic trace of the redox reaction between 1-(ferrocenyl)-ethanone/ethanol and dicyanobis(phenanthroline)iron(III) in 80% (v/v) aqueous dioxane at 530 nm, 0.08 mM [FeIII(phen)2(CN)2+], and 0.18 mM ‘I’. Plot “A” shows the effect of concentration of 1-ferrocenylethanone on the rate of reaction at 304 ± 0.5 K. Different colours represent the concentration range between 0.08 mM and 1.88 mM. Plot “B” displays the effect of 1-ferrocenylethanol on the rate of reaction at 291 ± 0.5 K. The concentration of 1-ferrocenylethanol was varied from 0.08 mM to 1.88 mM (coloured traces).
Figure 2 Kinetic trace of the redox reaction between 1-(ferrocenyl)-ethanone/ethanol and dicyanobis(phenanthroline)iron(III) in 80% (v/v) aqueous dioxane at 530 nm, 0.08 mM [FeIII(phen)2(CN)2+], and 0.18 mM ‘I’. Plot “A” shows the effect of concentration of 1-ferrocenylethanone on the rate of reaction at 304 ± 0.5 K. Different colours represent the concentration range between 0.08 mM and 1.88 mM. Plot “B” displays the effect of 1-ferrocenylethanol on the rate of reaction at 291 ± 0.5 K. The concentration of 1-ferrocenylethanol was varied from 0.08 mM to 1.88 mM (coloured traces).

A plot of absorbance versus time passes through origin and yields a slope (ε ⋅ b ⋅ kobs) according to the integrated rate equation of zero-order. The end Eq. (4) of the derivation (mentioned in supplementary information) is mentioned below. The symbols ε, b, and kobs have been used for the molar absorptivity (M−1 cm−1) of [FeII(phen)2(CN)2], the path length (cm) of the cuvette, and the observed zeroth order rate constant (M s−1), respectively. The path length of the cuvette was 1 cm, and the ε term is a constant value. These two factors do not change the overall behaviour of the rate constant, kobs. We calculated the slope instead of calculating the rate constant separately. The ‘slope’ has dimension s−1. This is noteworthy, the first order ‘rate constant’ also carries a unit s−1.

(3)
[ A Y ] t = ε bk obs t
(4)
∵ b = 1 cm [ A Y ] t = ε k obs t

Meanwhile, a plot of ln absorbance at times t (y-coordinate) and time (x-coordinate) yields a slope, which is equivalent to the observed pseudo-first order rate constant, k′obs. The intercept represents the value of ln absorbance at zero time (equation mentioned below). We observed an intercept in our plot, although we followed formation of [FeII(phen)2(CN)2] along time (derivation is mentioned in supplementary information). This system is normally supposed to have no concentration at zero time. In our studies, the reaction(s) was zeroth order in the first phase which rose [FeII(phen)2(CN)2]. This quantity was present in the reaction mixture at the start of the second phase with an overall second order kinetics.

(5)
ln [ A Y ] t = ln [ A Y ] 0 + k obs ′ t

This is also significant to mention here that the first-order plot was drawn to the best-fit data before it bent. This is because the third phase; phase of competition, made the measurements flawed near completion of the reaction(s). The plot(s) retained linearity with a reduced value(s) of the slope(s) after bent, which supports our assumption.

3.1

3.1 Effect of the concentration of reactants on the rate constants

Increasing concentration of [FeIII(phen)2(CN)2]+ did not affect the observed zeroth order (kobs) and pseudo-first order (k′obs) rate constants (Fig. 3). Together, these findings support the suitability of conditions that we maintained for obtaining the pseudo-first order kinetic data.

Effect of variation in the concentration of [FeIII(phen)2(CN)2]+ on ε kobs ( ), and k′obs ( ) when [CpFeIICpCOMe] or [CpFeIICpCHOHMe] was constant, 1.3 mM, and ionic strength was 0.18 mM. Impact of the concentration of [CpFeIICpCOMe] or [CpFeIICpCHOHMe] on the rate constants (ε kobs; , k′obs; ), respectively, at 0.08 mM [FeIII(phen)2(CN)2]+ and 0.18 mM ‘I’. The filled circle ( ) mentions k′obs at 1:1 concentration ratio between oxidant and reductant.
Figure 3 Effect of variation in the concentration of [FeIII(phen)2(CN)2]+ on ε kobs ( ), and k′obs ( ) when [CpFeIICpCOMe] or [CpFeIICpCHOHMe] was constant, 1.3 mM, and ionic strength was 0.18 mM. Impact of the concentration of [CpFeIICpCOMe] or [CpFeIICpCHOHMe] on the rate constants (ε kobs; , k′obs; ), respectively, at 0.08 mM [FeIII(phen)2(CN)2]+ and 0.18 mM ‘I’. The filled circle ( ) mentions k′obs at 1:1 concentration ratio between oxidant and reductant.

No considerable change was observed in the value of ε kobs upon increasing the concentration of [CpFeIICpCOMe] or [CpFeIICpCHOHMe] in the reaction mixture, keeping all other parameters constant. Contrary to that the value of k′obs increased. We found increasing curvature(s) in the graphs drawn between k′obs and [CpFeIICpCOMe] or [CpFeIICpCHOHMe]. The plots between [FeIII(phen)2(CN)2]+ (primary x-axis) and ε kobs (left, primary y-axis), and [CpFeIICpCOMe]/[CpFeIICpCHOHMe] (secondary x-axis) and k′obs (right, secondary y-axis) refer to our results (Fig. 3). The elevation pattern of each curve is different, where plot ‘A’ outlines 1-ferrocenylethanone and plot ‘B’ 1-ferrocenylethanol. A linear plot (Fig. 3B) demonstrates that the reaction is first order with respect to [CpFeIICpCHOHMe]. However, the curve with respect to [CpFeIICpCOMe] (Fig. 3A) stoops at the higher concentration of [CpFeIICpCOMe]. This indicated formation of other species of [CpFeIICpCOMe] in the reaction mixture, which decreased the concentration of [CpFeIICpCOMe], and affected the rate constant (k′obs), consequently. Protonation is a well-known tendency of 1-ferrocenylethanone (Arnett and Bushick, 1962; Khattak, 2011; Rubalcava and Thomson, 1963). We assume formation of the conjugate acid of 1-ferrocenylethanone; [CpFeIICpC+OHMe] in aqueous dioxane, and that its oxidation to FeIII state during the reaction. The oxygen atom of the carbonyl group in acetyl furnishes a site to protonation. The intercept of the curves (Fig. 3) confirms initial zeroth order phase of the reaction(s). A discordant value of k′obs at 1:1 concentration ratio between reducing and oxidising agents testifies our results, and also testifies initial zeroth order kinetics (Fig. 3).

Hydrogen bond exists between hydrogen and oxygen of water and dioxane in aqueous dioxane, which reduces pH of the solvent mixture. This provides a base to protonate the carbonyl group of 1-(ferrocenyl)-ethanone/ethanol. Autoionization or autoprotolysis occurs in co-solvent mixtures. The pH measurements need to include ‘correction factor’, consequently (Khattak, 2011). The revisited pH of 90, 80, 70, 60, and 50 (%, v/v) aqueous dioxane mixtures was found to be 3.96, 3.57, 3.57, 3.67, and 3.93 at 32 °C, respectively. The pH of 0.75 mM [CpFeIICpCOMe]/[CpFeIICpCHOHMe] in 80% (v/v) aqueous dioxane was calculated 3.69/3.75 at 32 °C. The pH of the water was 7.17 at 32 °C. These findings outline appropriateness of the condition to protonate 1-(ferrocenyl)-ethanone/ethanol.

3.2

3.2 Protonation and the rate constants

Formation of protonated 1-(ferrocenyl)-ethanone/ethanol and their reaction(s) to reduce dicyanobis(phenanthroline)iron(III) were examined. The oxygen atom of 1-ferrocenylethanol also maintains a site to interact with proton and generate [CpFeIICpCHO+H2Me]. The reaction(s) was probed at varying concentrations of protons keeping all other parameters constant. The concentration of 1-(ferrocenyl)-ethanone/ethanol was almost 10-fold excess over dicyanobis(phenanthroline)iron(III). Nitric acid was used to maintain the concentration of protons. The results showed a constant value of ε kobs (Fig. 4). The value of k′obs increased during oxidation of 1-ferrocenylethanone as the concentration of protons increased, and reduced in case of 1-ferrocenylethanol. A plot of [H+] versus ε kobs and k′obs summarised these profiles (Fig. 4). The results also outlined bending curvature at higher concentration of protons, which demonstrate [CpFeIICpCOMe] and [CpFeIICpCHOHMe] to be the limiting reactants. These outcomes also substantiate the bending curvature, Fig. 3A, when protons participated as the limiting reactant.

Reaction profile emerges from the effect of increasing concentration of protons on the rate constants (ε kobs; , k′obs; ) at 5.1 mM (I), 0.08 mM ([FeIII(phen)2(CN)2]+) and 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]).
Figure 4 Reaction profile emerges from the effect of increasing concentration of protons on the rate constants (ε kobs; , k′obs; ) at 5.1 mM (I), 0.08 mM ([FeIII(phen)2(CN)2]+) and 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]).

The results confirm protonation of each; 1-ferrocenylethanone and 1-ferrocenylethanol, and that, [CpFeIICpC+OHMe] participates in the rate-determining step whilst [CpFeIICpCOMe] does not affect the rate constant i.e., zeroth order phase of reaction (Fig. 4A). Protonation of 1-ferrocenylethanol reduced the concentration of free 1-ferrocenylethanol in the reaction mixture. The value of pseudo-first order rate constant decreased consequently upon increasing the concentration of protons in the reaction mixture. These outcomes corroborate [CpFeIICpCHOHMe] as an active species leading the rate-determining step, and zero-order with respect to [CpFeIICpCHO+H2Me] (Fig. 4B).

3.3

3.3 Active species and the rate-determining step(s)

We tried to wet our results in the view of transition state theory. Thus the primary salt effect and effect of dielectric constant were monitored on the rate constants in order to single out the active species, which contributes to the rate-determining step(s). The ionic strength was varied by potassium nitrate, and the dielectric constant was maintained through changing proportion of 1,4-dioxane to water i.e., % v/v. The zeroth order data (ε kobs) were unaffected by either factor whether primary salt effect or dielectric constant (Figs. 5 and 6). Contrary to that there are variations in k′obs. Instead of the value of the slope (2A zA zB) of the plot of log k′obs versus √I being zero, as one of the reactant is neutral, we found a nonzero slope in case of [CpFeIICpCOMe] and a zero slope corresponding to [CpFeIICpCHOHMe] (Fig. 5). The results confirmed oxidation of [CpFeIICpCHOHMe] during the rate-determining step and [CpFeIICpCHO+H2Me] in the first phase of the reaction (zero-order phase). Scaling of ionic strength increased the value of k′obs when [CpFeIICpCOMe] reacted with [FeIII(phen)2(CN)2]+. We obtained a linear graph between log k′obs and √I (Fig. 5A). This clearly suggests the active species carried same charges (+ve) such as [CpFeIICpC+OHMe] and [FeIII(phen)2(CN)2]+, and also that, the rate-determining step involved these species. We observed same results in a range of co-solvent mixtures (aqueous dioxane). The plots were drawn excluding out layers, which yielded slope amid 15–28. The values are although high, but we may consider them normal. This is because Debye–Hückel constant (A) depends upon dielectric constant, nature of solvent, and temperature. Another probability may be a low range of ionic strength we used bearing in mind the proportion of dioxane (90–50%, v/v aqueous dioxane) to avoid precipitation of salt. The intercept of the plot of log k′obs versus √I, yields log k′obs in ideal condition. The condition is considered ideal if ionic strength is zero. We studied the effect of dielectric constant on the rate constants under ideal conditions in order to avoid any interference by ionic strength. The natural logarithm of ideal k′obs was plotted against the reciprocal of the dielectric constant (Fig. 6A) according to formulation mentioned below. The plot produced a negative slope, which confirmed the active species are positively charged such as [CpFeIICpC+OHMe] and [FeIII(phen)2(CN)2]+. We figured out the inter-nuclear distance (r#) between the charges in the activated complex from the slope, and found 1.38 nm. The change in the dielectric constant did not produce any alteration in the values of ε kobs and k′obs when the reactants were [CpFeIICpCHOHMe] and [FeIII(phen)2(CN)2]+ (Fig. 6B). These results confirmed once again that the rate-determining step involved a neutral reactant such as [CpFeIICpCHOHMe] which produces a zero slope.

(6)
ln k obs ′ ideal = ln k 0 - z A z B e 2 4 π ε 0 ε r r ≠ kT
The impact of variation in the ionic strength on k′obs (secondary x–y coordinates) and ε kobs (primary x–y coordinates) at 0.08 mM ([FeIII(phen)2(CN)2]+) and 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]).
Figure 5 The impact of variation in the ionic strength on k′obs (secondary x–y coordinates) and ε kobs (primary x–y coordinates) at 0.08 mM ([FeIII(phen)2(CN)2]+) and 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]).
Effect of dielectric constant of the medium on the rate constants, k′obs ( ) and ε kobs ( ). The experimental conditions were maintained at 0.08 mM ([FeIII(phen)2(CN)2]+), 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]) and zero ionic strength.
Figure 6 Effect of dielectric constant of the medium on the rate constants, k′obs ( ) and ε kobs ( ). The experimental conditions were maintained at 0.08 mM ([FeIII(phen)2(CN)2]+), 0.75 mM ([CpFeIICpCOMe] or [CpFeIICpCHOHMe]) and zero ionic strength.

4

4 Proposed mechanism(s)

We thus propose the mechanism of each reaction in view of this study.

4.1

4.1 Oxidation of 1-ferrocenylethanone by dicyanobis(phenanthroline)iron(III)

This study revealed equilibrium between [CpFeIICpCOMe] and [CpFeIICpC+OHMe] in aqueous dioxane. Protonated and unprotonated, 1-ferrocenylethanone reduced [FeIII(phen)2(CN)2]+. The reaction was first order with respect to [FeIII(phen)2(CN)2]+, and [CpFeIICpC+OHMe], i.e., an overall second order reaction. This phase (second phase) of the reaction was the rate-determining or slow-step. The first phase consisted of the reaction between [CpFeIICpCOMe] and [FeIII(phen)2(CN)2]+, when the reaction followed zero-order kinetics. Following are the stepwise reactions.

(7)
(8)
[ Fe III ( phen ) 2 ( CN ) 2 ] + + [ CpFe II CpCOMe ] → k 1 [ Fe II ( phen ) 2 ( CN ) 2 ] + [ CpFe III CpCOMe ] +
(9)
[ Fe III ( phen ) 2 ( CN ) 2 ] + + [ CpFe II CpC + OHMe ] → k 2 [ Fe II ( phen ) 2 ( CN ) 2 ] + [ CpFe III CpCOMe ] + + [ H ] +

The rate of formation of [FeII(phen)2(CN)2] is as follows:

(10)
d [ Fe II ( phen ) 2 ( CN ) 2 ] dt = k 1 [ Fe III ( phen ) 2 ( CN ) 2 + ] 0 [ CpFe II CpCOMe ] 0 + k 2 [ Fe III ( phen ) 2 ( CN ) 2 + ] [ CpFe II CpC + OHMe ]

Eq. (10) becomes reduced to Eq. (11),

(11)
d [ Fe II ( phen ) 2 ( CN ) 2 ] dt = k 1 + k 2 [ Fe III ( phen ) 2 ( CN ) 2 + ] [ CpFe II CpC + OHMe ]

The total concentration of 1-ferrocenylethanone equals to the addition product of [CpFeIICpC+OHMe] and [CpFeIICpCOMe]F, and an equilibrium exists between [CpFeIICpC+OHMe] and [CpFeIICpCOMe]F.

(12)
[ CpFe II CpCOMe ] T = [ CpFe II CpC + OHMe ] + [ CpFe II CpCOMe ] F
(13)
K eq = [ CpFe II CpC + OHMe ] [ CpFe II CpCOMe ] F [ H + ]

The concentration of [CpFeIICpC+OHMe] is obtained if we put value of [CpFeIICpCOMe]F in Eq. (12) from Eq. (13) and solve. This resulting value upon insertion in Eq. (11) gives the following:

(14)
[ CpFe II CpC + OHMe ] = K eq [ CpFe II CpCOMe ] T [ H + ] 1 + K eq [ H + ]
(15)
d [ Fe II ( phen ) 2 ( CN ) 2 ] dt = k 1 + k 2 K eq [ Fe III ( phen ) 2 ( CN ) 2 + ] [ CpFe II CpCOMe ] T [ H + ] 1 + K eq [ H + ]

We studied the reaction under pseudo-first order condition. The concentration of [CpFeIICpCOMe] was always in several folds excess over [FeIII(phen)2(CN)2]+. Under these conditions, the rate was dependent upon the concentration of [FeIII(phen)2(CN)2]+, when the proportionality constant was equal to the observed rate constant. The observed rate constant was proportional to [CpFeIICpCOMe], which produced proportionality constant. This constant was equal to the overall rate constant. Added to this, the concentration term [FeIII(phen)2(CN)2+] yielded equivalent concentration of [FeII(phen)2(CN)2] bearing in mind the one electron-exchange chemistry. We obtained Eqs. (16) and (17) upon dividing Eq. (15) by [FeII(phen)2(CN)2].

(16)
k obs ′ = k obs + k 2 K eq [ CpFe II CpCOMe ] T [ H + ] 1 + K eq [ H + ]
(17)
1 k obs ′ = 1 k obs + 1 + K eq [ H + ] k 2 K eq [ CpFe II CpCOMe ] T [ H + ]

Eq. (16) demonstrates that the observed pseudo-first order rate constant (k′obs) depends upon the first power of the concentration of 1-ferrocenylethanone (total), and protons, whilst the value of kobs is constant. Our results are in good agreement with the proposed rate law.

An approximate value of Keq (1.59 mM−1) was calculated by taking reciprocal of the acid dissociation constant (Ka) of [CpFeIICpC+OHMe]. Earlier studies reported the value of Ka, 631 M (pKa = −2.80 in aqueous sulphuric acid solution) (Arnett and Bushick, 1962; Khattak, 2011; Rubalcava and Thomson, 1963). However, the value of Ka may probably be little different from the reported one in 80% (v/v) aqueous dioxane. The multiplication product of Keq and [H+] (∼0.0001–0.001 M in the reaction mixture) reduced (1 + Keq [H+]) to 1. Eq. (17) reduced to Eq. (18), consequently.

(18)
1 k obs ′ = 1 k obs + 1 k 2 K eq [ CpFe II CpCOMe ] T [ H + ]

Eq. (18) assigns a linear relation to the reciprocals of k′obs and concentration of 1-ferrocenylethanone, or protons (Fig. 7). The intercept(s) of the plot(s) yielded kobs, which produced k1 (an overall zero-order rate constant). We figured out k2 (an overall second order rate constant) from the slope(s) of the plot(s). An average value of k1 is 0.42 mM s−1. We calculated two different numerals of k2. Among the two values, 6.19 × 1010 M−1 s−1 and 1.03 × 1011 M−1 s−1, we suppose 1.03 × 1011 M−1 s−1 comparatively precise, because we used a true value of [CpFeIICpCOMe]T in calculations rather than an estimated figure of [H+], which was determined from pH.

Linear relationship between the reciprocals of k′obs, [CpFeIICpCOMe]T, and [H+]; (1/k′obs − 1/[CpFeIICpCOMe]T, ) and (1/k′obs − 1/[H+], ).
Figure 7 Linear relationship between the reciprocals of k′obs, [CpFeIICpCOMe]T, and [H+]; (1/k′obs − 1/[CpFeIICpCOMe]T, ) and (1/k′obs − 1/[H+], ).

4.2

4.2 Oxidation of 1-ferrocenylethanol by dicyanobis(phenanthroline)iron(III)

Our findings show equilibrium between [CpFeIICpCHOHMe] and [CpFeIICpCHO+H2Me] in aqueous dioxane. Initially, in the first phase, the reaction involved electron transfer between [CpFeIICpCHO+H2Me] and [FeIII(phen)2(CN)2]+. The reaction was zero-order in this phase. An overall second order reaction phase was proved to be the rate-determining or slow-step. The reaction was first order with respect to [FeIII(phen)2(CN)2]+ and [CpFeIICpCHOHMe] in this phase. The step-wise reactions are as follows:

(19)
(20)
[ Fe III ( phen ) 2 ( CN ) 2 ] + + [ CpFe II CpCHO + H 2 Me ] → k 1 [ Fe II ( phen ) 2 ( CN ) 2 ] + [ CpFe III CpCHOHMe ] + + [ H ] +
(21)
[ Fe III ( phen ) 2 ( CN ) 2 ] + + [ CpFe II CpCHOHMe ] → k 2 [ Fe II ( phen ) 2 ( CN ) 2 ] + [ CpFe III CpCHOHMe ] +

The rate of formation of dicyanobis(phenanthroline)iron(II) is as follows,

(22)
d [ Fe II ( phen ) 2 ( CN ) 2 ] dt = k 1 [ Fe III ( phen ) 2 ( CN ) 2 + ] 0 [ CpFe II CpCHO + H 2 Me ] 0 + k 2 [ Fe III ( phen ) 2 ( CN ) 2 + ] [ CpFe II CpCHOHMe ] F

Eq. (22) reduces to Eq. (23), if we observe algorithmic rules.

(23)
d [ Fe II ( phen ) 2 ( CN ) 2 ] dt = k 1 + k 2 [ Fe III ( phen ) 2 ( CN ) 2 + ] [ CpFe II CpCHOHMe ] F

We know, [CpFeIICpCHOHMe]T is equal to [CpFeIICpCHO+H2Me] and [CpFeIICpCHOHMe]F, and that both species are in equilibrium.

(24)
[ CpFe II CpCHOHMe ] T = [ CpFe II CpCHOHMe ] F + [ CpFe II CpCHO + H 2 Me ]
(25)
[ CpFe II CpCHO + H 2 Me ] = K eq [ CpFe II CpCHOHMe ] F [ H + ]

If we insert the value of [CpFeIICpCHO+H2Me] from Eq. (25) in Eq. (24), we obtained Eq. (26).

(26)
[ CpFe II CpCHOHMe ] F = [ CpFe II CpCHOHMe ] T 1 + K eq [ H + ]

We get Eq. (27) by inserting the value of [CpFeIICpCHOHMe]F in Eq. (23) bearing in mind the pseudo-first order condition.

(27)
k obs ′ = k obs + k 2 [ CpFe II CpCHOHMe ] T 1 + K eq [ H + ]

Eq. (27) solves to Eq. (28), if we consider the effect of protons.

(28)
1 k obs ′ = 1 k obs + 1 k 2 [ CpFe II CpCHOHMe ] T + K eq [ H + ] k 2 [ CpFe II CpCHOHMe ] T ∵ 1 k obs + 1 k 2 [ CpFe II CpCHOHMe ] T = intercept

A linear graph must obtain when plotted between 1/k′obs and [H+]. Our results agreed with the proposed mechanism (Fig. 8). The slope of the plot gave Keq 300 M−1, where the value of [CpFeIICpCHOHMe]T and k2 were known (Fig. 3B). We calculated Ka and pKa of [CpFeIICpCHO+H2Me] by using Keq as follows: K a = 1 K eq = 1 300 M - 1 = 3.33 mM pK a = - log K a = 2.48

Linear relationship between 1/k′obs and [H+]. Inset shows the behaviour of [CpFeIICpCHOHMe] as a limiting reactant at the higher concentration of protons. The reaction parameters were set at 5.1 mM (I), 0.08 mM ([FeIII(phen)2(CN)2]+), 0.75 mM [CpFeIICpCHOHMe], and 291 ± 0.5 K.
Figure 8 Linear relationship between 1/k′obs and [H+]. Inset shows the behaviour of [CpFeIICpCHOHMe] as a limiting reactant at the higher concentration of protons. The reaction parameters were set at 5.1 mM (I), 0.08 mM ([FeIII(phen)2(CN)2]+), 0.75 mM [CpFeIICpCHOHMe], and 291 ± 0.5 K.

The term; (1 + Keq[H+]) in Eq. (27) equals to 1, if Keq (300 M−1) is multiplied to the concentration of protons (i.e. ∼10−4 M) present in the reaction mixture (80% aqueous dioxane).

(29)
k obs ′ = k obs + k 2 [ CpFe II CpCHOHMe ] T

Plot of k′obs versus [CpFeIICpCHOHMe]T (Fig. 3B) identifies the mechanism reliable, and produced k2 1.38 × 103 M−1 s−1.

5

5 Conclusion

This study explored the pivotal role of protonated 1-(ferrocenyl)-ethanone/ethanol during the reaction of 1-(ferrocenyl)-ethanone/ethanol. Dicyanobis(phenanthroline)iron(III) does not protonate under the experimental conditions. The relationship of Gibbs free energy, equilibrium constant, and cell potential, as well as the Marcus cross-relation, helped to determine the values of the electrode potential (E1/2) and the self-exchange rate constant (k22) of [CpFeIICpC+OHMe]. This is an unsaturated intermediate species, which gets formed in aqueous dioxane. Reported values of the electrode potential and self-exchange rate constants of dicyanobis(phenanthroline)iron(III/II) ([FeIII/II(phen)2(CN)2]+/0), 2-ferrocenylethanol ([CpFeIII/IICpCH2CH2OH]+/0) and 1-ferrocenylethanol ([CpFeIII/IICpCHOHMe]+/0) couples were used to figure out E1/2 (0.204 V) and k22 (5.7 × 104 M−1 s−1) of [CpFeIII/IICpC+OHMe]+/0 (Castillo, 1986; Garcia et al., 1987; Matsumoto et al., 1997; Ryabov et al., 1999; Takagi et al., 1995; Wynn et al., 1984). This is significant to consider the structural difference between 1-ferrocenylethanol, 2-ferrocenylethanol, and protonated 1-ferrocenylethanone. The former two metallocenes are isomers but their properties, reduction potential and self-exchange rate constant, are different due to the position of hydroxyl group. The hydroxyl group in 2-ferrocenylethanol does not strongly withdraw electrons as compared to 1-ferrocenylethanol. Protonated 1-ferrocenylethanone is different from 1-ferrocenylethanol in such a way that the latter is saturated at the carbon attached to oxygen, and is very stable. The solubility of this species is also fair in water as compared to 1-ferrocenylethanone, which is sparingly soluble. Protonated 1-ferrocenylethanone is unsaturated at carbon attached to oxygen to form a carbonium ion. This is a highly reactive species and participates in the rate-determining step. E1/2 (0.204 V) of [CpFeIII/IICpC+OHMe]+/0 is identical to E1/2 (0.204 V) of 1-ferrocenylethanol (pH 7, 25 °C). It is remarkable that the E1/2 of [CpFeIII/IICpC+OHMe]+/0 may probably be greater in value (>0.204 V) in the 80% (v/v) aqueous dioxane. Therefore, calculated E1/2 of [CpFeIII/IICpC+OHMe]+/0 is an approximate value. We also observed that 1-ferrocenylethanol participates in the rate-determining step during the reduction of dicyanobis(phenanthroline)iron(III). Together, these findings lead to a conclusion that the rate of oxidation of 1-ferrocenylethanone and 1-ferrocenylethanol is restricted to the protonation of carbonyl group either unsaturated at carbon centre; [CpFeIICpC+OHMe] or, saturated at ; [CpFeIICpCHOHMe]. The observed values of pseudo-first order rate constant for each reaction are almost similar. This confirms that protonation of the carbonyl group is an important factor in the oxidation of the derivatives of ferrocene. This may change the mechanisms slightly or completely and affix margin to the properties, chemical reactivity and biological effects of these compounds.

Acknowledgement

Authors acknowledge the services of H.E.J. Research Institute of Chemistry, University of Karachi, to analyse the synthesised compounds. Such services include elemental analysis. We are thankful to Dr. David R.W. Hodgson, Department of Chemistry, Durham University, UK, for some technical and linguistic assistance.

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Appendix A

Supplementary material

Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.arabjc.2016.05.007.

Appendix A

Supplementary material

Supplementary data 1

Supplementary data 1


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