# Antioxidative Reactivity of L-Ascorbic Acid and D-Isoascorbic Acid Species towards Reduction of Hexachloroiridate (IV)

Authors: Han Zhang; Yanqing Xia; Peng Zhang; Liqian Hou; Ying Sun; Zhaozhan Lu; Yunfeng Tang; Hongwu Tian; Tiesheng Shi
Journal: Journal of Chemistry
 (2021)

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## Abstract

The pair [IrCl6]2–/[IrCl6]3– has been demonstrated to be a good redox probe in biological systems while L-ascorbic acid (AA) is one of the most important antioxidants. D-isoascorbic acid (IAA) is an epimer of AA and is widely used as an antioxidant in various foods, beverages, meat, and fisher products. Reductions of [IrCl6]2– by AA and IAA have been analyzed kinetically and mechanistically in this work. The reductions strictly follow overall second-order kinetics and the observed second-order rate constants were collected in the pH region of 0 ≤ pH ≤ 2.33 at 25.0°C. Spectrophotometric titration experiments revealed a well-defined 1 : 2 stoichiometry, namely Δ[AA] : Δ[Ir(IV)] or Δ[IAA] : Δ[Ir(IV)] = 1 : 2, indicating that L-dehydroascorbic acid (DHA) and D-dehydroisoascorbic acid (DHIA) were the oxidation products of AA and IAA, respectively. A reaction mechanism is suggested involving parallel reactions of [IrCl6]2– with three protolysis species of AA/IAA (fully protonated, monoanionic, and dianionic forms) as the rate-determining steps and formation of ascorbic/isoascorbic and ascorbate/isoascorbate radicals; in each of the steps, [IrCl6]2– acquires an electron via an outer-sphere electron transfer mode. Rate constants of the rate-determining steps have been derived or estimated. The fully protonated forms of AA and IAA display virtually identical reactivity whereas ascorbate and isoascorbate monoanions have a significant reactivity difference. The ascorbate and isoascorbate dianions are extremely reactive and their reactions with [IrCl6]2– proceed with the diffusion-controlled rate. The species versus pH and the species reactivity versus pH distribution diagrams were constructed endowing that the ascorbate/isoascorbate monoanionic form dominated the total reactivity at physiological pH. In addition, the value of pKa1 = 3.74 ± 0.05 for IAA at 25.0°C and 1.0 M ionic strength was determined in this work.


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## Body

## 1. Introduction

Vitamin C is a general name for all the compounds which possess the fully or partially biological activity of L-ascorbic acid (AA) [1–6]. It includes AA itself, L-dehydroascorbic acid (DHA), the esters of AA such as ascorbyl palmitate, and D-isoascorbic acid (IAA); the latter two are the synthetic compounds [1–3]. Vitamin C is widely found in plants (vegetables and fruits in particular) and in animals, and nowadays it is also available as caplets, tablets, capsules, and drink mixes and in multivitamin and antioxidant formulations [1–3]. Biologically and/or medically, extensive investigations of vitamin C have been carried out [7, 8]. When vitamin C is utilized in food and nutrients, its antioxidative property is often considered to be an important factor [9, 10]. IAA is an epimer of AA and widely used as an antioxidant in various foods, beverages, meat, and fisher products [2, 11–15]. In this regard, several analytical methods have been developed to determine content of IAA in foodstuffs, together with those of AA and their oxidized forms [2, 11–15]. Scheme 1 provides the structures of AA, IAA and their oxidized forms.Scheme 1
Structures of L-ascorbic acid and D-isoascorbic acid and of their common oxidation forms L-dehydroascorbic acid and D-dehydroisoascorbic acid.As one of the most important antioxidants, the kinetics and mechanisms of vitamin C/ascorbic acid reactions with various oxidants have been explored extensively [9, 10, 16–20]; the oxidants include reactive oxygen species such as HO2/O2– radicals and peroxynitrite, and some metal ions/complexes [16–19]. Among the metal ions/complexes, single electron oxidants apparently dominate the exploration [9, 10]. We have had an interest in the study of the redox pair [IrCl6]2–/[IrCl6]3– due to its salient characters: (i) [IrCl6]2– has two strong absorption bands (around 420 and 488 nm, vide infra) and [IrCl6]3– does not possess these two bands, conferring a convenient monitoring of the pair. (ii) [IrCl6]2– is very stable and meta-stable from pH 0 to 10, enabling us to measure the rate constants in such a wide pH range [20–22]. (iii) It has been used as a redox probe to discriminate DNA monolayers and to obtain chemical information of oxidative stress in biological matrixes [23, 24].The antioxidative reactivity of AA and IAA towards reduction of [IrCl6]2– has been investigated in this work. Previously, the kinetics of oxidation of AA by [IrCl6]2– was studied [25, 26], but only was carried out in a very narrow acidity region (namely, 0.20 M ≤ [HClO4] ≤ 1.00 M). On the other hand, the redox reaction between [IrCl6]2– and IAA has never been investigated. In fact, the oxidation kinetic and mechanistic aspects of the IAA oxidations were poorly understood due to a very limited study [19]. The purposes of the present work were to study the redox reactions in a pH range as wide as possible, to derive the rate constants of the rate-determining steps in a good accuracy, to propose a convincing reaction mechanism, and to find out the antioxidative difference between AA and IAA.


## 2. Experimental Section

### 2.1. Materials

Sodium hexachloroiridate(IV) hexahydrate (Na2IrCl6·6H2O, 99.9%) and L-ascorbic acid (AA, 99%) were purchased from Sigma-Aldrich. D-(-)-isoascorbic acid (IAA, HPLC grade, purity ≥99%) was obtained in its purest form from Yuanye Biological Company (Shanghai, China). Concentrated perchloric acid, sodium perchlorate monohydrate, concentrated o-phosphoric acid (ca. 85%), sodium dihydrogen-phosphate, disodium hydrogen phosphate, acetic acid, sodium acetate, sodium chloride, and 0.100 M NaOH standard solution were purchased in the purest forms available either from Fisher Scientific (Fisher Scientific, Shanghai, China) or from Alfa Aesar (Alfa Aesar Branch in Shanghai, China). These chemicals were used directly without further purification. For pH meter calibrations, standard buffers of pH 4.00, 7.00, and 10.00 were purchased from ThermoFisher. Doubly distilled water was used throughout this work.


### 2.2. Reaction Media or Buffer Solutions

Stock solutions of 1.00 M HClO4 and H3PO4 were prepared from the concentrated acids and standardized by the 0.100 M NaOH solution. For pH ≤ 2.11, the reaction media were prepared by combinations of the 1.00 M HClO4 with 1.00 M NaClO4; their pH values were derived by use of the equation: pH = -log[H+] + 0.11, which is based on a mean activity coefficient of 0.77 in solutions of 1.00 M NaClO4 [27]. For buffer solutions in the pH range from 2.28 to 6.26, the combinations of the buffering pairs H3PO4/NaH2PO4, AcOH/NaOAc, and NaH2PO4/Na2HPO4 (0.15–0.2 M) were employed. All the buffers were tuned to an ionic strength of 1.0 M by use of NaClO4·H2O; their pH values were measured by an Accumet Basic AB150 Plus pH meter, equipped with an Accumet accuTupH® combination pH electrode (ThermoFisher). The electrode was calibrated using standard buffers of pH 4.00, 7.00, and 10.00 immediately before pH measurements. These buffer solutions were only used weekly.


### 2.3. Kinetic Experiments

Stock solutions of about 10 mM [IrCl6]2– were prepared daily by dissolving the desired amount of Na2IrCl6·6H2O in a solution containing 0.90 M NaClO4, 0.09 M NaCl, and 0.01 M HCl; these solutions were only used daily. Stock solutions of AA/IAA were prepared by adding the desired amount of AA/IAA to a buffer solution of specific pH. Each of the solutions was bubbled with pure nitrogen for 5 min and was only employed for ca. 2 h. Solutions of [IrCl6]2– and AA/IAA for kinetic measurements were made by dilution of the above stock solutions with the same pH buffer; these solutions were loaded onto an Applied Photophysics SX-20 stopped-flow spectrometer (Applied Photophysics Ltd., Leatherhead, UK). A water bath circulation from a thermostat (Lauda Alpha RA8, Delran, NJ, USA) was employed for the temperature control at 25.0 ± 0.1°C. Kinetic traces were recorded by mixing equal volumes of the [IrCl6]2– and AA/IAA solutions directly in the stopped-flow spectrometer. Pseudo-first-order conditions (namely, [AA]tot or [IAA]tot ≥ 10∙[Ir(IV)]) were always maintained, where [AA]tot and [IAA]tot are ascribed to the total concentrations of AA and IAA, respectively.


### 2.4. Spectrophotometric Titration Experiments

UV-Vis spectra were recorded on a TU-1900 spectrophotometer by use of 1.00 cm quartz cells at 25.0°C (Beijing Persee, Beijing, China); the temperature was controlled by a water circulation from the same type of thermostat. There were series of reaction mixtures in which [Ir(IV)] = 0.40 mM was kept unchanged and [AA]tot/[IAA]tot was varied from 0 to 0.50 mM in a reaction medium of [H+] = 0.10 M and 1.0 M ionic strength. Each of the reaction mixtures was controlled at 25°C for about 5 min; the absorption value was measured at 488 nm.


## 2.1. Materials

Sodium hexachloroiridate(IV) hexahydrate (Na2IrCl6·6H2O, 99.9%) and L-ascorbic acid (AA, 99%) were purchased from Sigma-Aldrich. D-(-)-isoascorbic acid (IAA, HPLC grade, purity ≥99%) was obtained in its purest form from Yuanye Biological Company (Shanghai, China). Concentrated perchloric acid, sodium perchlorate monohydrate, concentrated o-phosphoric acid (ca. 85%), sodium dihydrogen-phosphate, disodium hydrogen phosphate, acetic acid, sodium acetate, sodium chloride, and 0.100 M NaOH standard solution were purchased in the purest forms available either from Fisher Scientific (Fisher Scientific, Shanghai, China) or from Alfa Aesar (Alfa Aesar Branch in Shanghai, China). These chemicals were used directly without further purification. For pH meter calibrations, standard buffers of pH 4.00, 7.00, and 10.00 were purchased from ThermoFisher. Doubly distilled water was used throughout this work.


## 2.2. Reaction Media or Buffer Solutions

Stock solutions of 1.00 M HClO4 and H3PO4 were prepared from the concentrated acids and standardized by the 0.100 M NaOH solution. For pH ≤ 2.11, the reaction media were prepared by combinations of the 1.00 M HClO4 with 1.00 M NaClO4; their pH values were derived by use of the equation: pH = -log[H+] + 0.11, which is based on a mean activity coefficient of 0.77 in solutions of 1.00 M NaClO4 [27]. For buffer solutions in the pH range from 2.28 to 6.26, the combinations of the buffering pairs H3PO4/NaH2PO4, AcOH/NaOAc, and NaH2PO4/Na2HPO4 (0.15–0.2 M) were employed. All the buffers were tuned to an ionic strength of 1.0 M by use of NaClO4·H2O; their pH values were measured by an Accumet Basic AB150 Plus pH meter, equipped with an Accumet accuTupH® combination pH electrode (ThermoFisher). The electrode was calibrated using standard buffers of pH 4.00, 7.00, and 10.00 immediately before pH measurements. These buffer solutions were only used weekly.


## 2.3. Kinetic Experiments

Stock solutions of about 10 mM [IrCl6]2– were prepared daily by dissolving the desired amount of Na2IrCl6·6H2O in a solution containing 0.90 M NaClO4, 0.09 M NaCl, and 0.01 M HCl; these solutions were only used daily. Stock solutions of AA/IAA were prepared by adding the desired amount of AA/IAA to a buffer solution of specific pH. Each of the solutions was bubbled with pure nitrogen for 5 min and was only employed for ca. 2 h. Solutions of [IrCl6]2– and AA/IAA for kinetic measurements were made by dilution of the above stock solutions with the same pH buffer; these solutions were loaded onto an Applied Photophysics SX-20 stopped-flow spectrometer (Applied Photophysics Ltd., Leatherhead, UK). A water bath circulation from a thermostat (Lauda Alpha RA8, Delran, NJ, USA) was employed for the temperature control at 25.0 ± 0.1°C. Kinetic traces were recorded by mixing equal volumes of the [IrCl6]2– and AA/IAA solutions directly in the stopped-flow spectrometer. Pseudo-first-order conditions (namely, [AA]tot or [IAA]tot ≥ 10∙[Ir(IV)]) were always maintained, where [AA]tot and [IAA]tot are ascribed to the total concentrations of AA and IAA, respectively.


## 2.4. Spectrophotometric Titration Experiments

UV-Vis spectra were recorded on a TU-1900 spectrophotometer by use of 1.00 cm quartz cells at 25.0°C (Beijing Persee, Beijing, China); the temperature was controlled by a water circulation from the same type of thermostat. There were series of reaction mixtures in which [Ir(IV)] = 0.40 mM was kept unchanged and [AA]tot/[IAA]tot was varied from 0 to 0.50 mM in a reaction medium of [H+] = 0.10 M and 1.0 M ionic strength. Each of the reaction mixtures was controlled at 25°C for about 5 min; the absorption value was measured at 488 nm.


## 3. Results and Discussion

### 3.1. Spectral Insight into the Reaction Courses

Although an outer-sphere electron transfer was suggested for the oxidation of AA by [IrCl6]2–, the reaction course was not exploited previously [25, 26]. The rapid scan spectra were thus recorded for the oxidations of AA/IAA by [IrCl6]2– in this work; Figure 1 displays such spectra in the case of AA. Clearly the absorption bands around 488, 420, and 306 nm, which are the typical ones of [IrCl6]2–, remained upshifted; in addition, no new absorption bands emerged during the reaction course. Moreover, the kinetic traces recorded at these bands could be well described by single exponentials, giving rise to the values of pseudo-first-order constants kobsd. The above reaction characters indicate that a simple electron transfer took place without complications such as an adduct formation or faster substitution reactions on [IrCl6]2– before the rate-determining step(s). It is thus concluded that the redox reaction is indeed first-order in [Ir(IV)]. The rapid scan spectra in the case of IAA are very similar to those shown in Figure 1.Figure 1
Rapid scan spectra for reduction of [IrCl6]2– by L-ascorbic acid (AA). Reaction conditions: [Ir(IV)] = 0.20 mM, [AA]tot = 2.00 mM, [H+] = 0.200 M, μ = 1.0 M, temperature of 25.0°C. The spectra were acquired at 2, 5, 10, 15, 20, 25, 30, 35, 40, 45, 50, 60, 70, and 80 milliseconds (ms) after mixing.


### 3.2. Second-Order Kinetics

Still under pseudo-first-order conditions, the reaction rates were determined when [AA]tot/[IAA]tot was varied in the region 0.10 ≤ [AA]tot/[IAA]tot ≤ 2.0 mM; for this concentration region, the reaction media had enough buffering capacities to ascertain that no pH changes could be caused by the variations of the concentrations of the reductants. The values of kobsd as functions of [AA]tot or [IAA]tot and of pH were collected; for each concentration, the kobsd value is reported as an average of the at least 5 duplicate runs and the standard deviations were usually less than 5%. Plots of kobsd versus [AA]tot are displayed in Figure 2 and versus [IAA]tot are shown in Figure 3; undoubtedly the linearity of the plots is very good, and no significant intercepts are observable. Hence, the redox reactions are first-order in [AA]tot or [IAA]tot, rendering overall second-order kinetics as described by equations (1a) and (1b), where k′ stands for the observed second-order rate constants.(1a)−dIrIVdt=k′AAtotIrIV,(1b)−dIrIVdt=k′IAAtotIrIV.Figure 2
Dependencies ofkobsd on [AA]tot at 25.0°C and 1.0 M ionic strength for reduction of [IrCl6]2– by AA.Figure 3
Dependencies ofkobsd on [IAA]tot at 25.0°C and 1.0 M ionic strength for reduction of [IrCl6]2– by IAA.Values ofk′ were calculated from the plots in Figures 2 and 3 and are summarized in Table 1. In addition, plots of logk′ versus pH are shown in Figure 4 (data points). The stopped-flow spectrometer was employed essentially to its up limit for following the redox reactions.Table 1
Observed second-order rate constants for the reduction of [IrCl6]2– by AA and IAA at 25.0°C and 1.0 M ionic strength.
ReductantpHk′/M−1s−1AA0.16(7.6 ± 0.1) × 1030.26(9.6 ± 0.2) × 1030.41(1.66 ± 0.02) × 1040.63(2.42 ± 0.04) × 1040.81(3.01 ± 0.06) × 1041.11(5.49 ± 0.09) × 1041.41(1.29 ± 0.03) × 1051.63(2.19 ± 0.05) × 1051.81(3.75 ± 0.09) × 1052.11(6.8 ± 0.2) × 1052.33(1.15 ± 0.06) × 106IAA0.16(7.5 ± 0.1) × 1030.26(8.3 ± 0.2) × 1030.41(1.40 ± 0.02) × 1040.63(2.04 ± 0.04) × 1040.81(2.78 ± 0.06) × 1041.11(4.78 ± 0.08) × 1041.41(9.4 ± 0.03) × 1041.63(1.97 ± 0.04) × 1051.81(3.10 ± 0.07) × 1052.11(6.1 ± 0.1) × 1052.33(1.00 ± 0.02) × 106Figure 4
Second-order rate constantsk′ (in logarithmic scale) versus pH plots at 25.0°C and 1.0 M ionic strength (data points) for reduction of [IrCl6]2– by AA and IAA. The solid lines represent the best fits of equation (6) to the experimental data by use of a weighted nonlinear least-squares method.


### 3.3. The Reaction Stoichiometry

The stoichiometric ratios were determined in a medium of 0.10 M HClO4 and 0.90 M NaClO4 by use of spectrophotometric titration method which was demonstrated to be a very useful approach [28–31]. The absorption values at 488 nm as a function of [AA]tot or [IAA]tot are shown in Figure 5; in either case, the data points all follow two crossing straight lines. The intersection points in Figure 5 confer stoichiometric ratios: Δ[Ir(IV)] : Δ[AA]tot = 0.40 mM : 0.202 mM = 1 : (0.51 ± 0.01) and Δ[Ir(IV)] : Δ[IAA]tot = 0.40 mM : 0.196 mM = 1 : (0.49 ± 0.01). Thus, the derived stoichiometric ratios Δ[AA]tot : Δ[Ir(IV)] and Δ[IAA]tot : Δ[Ir(IV)] = 1 : 2 are established unambiguously when the experimental errors are considered. This is perhaps not surprising because commonly AA is oxidized to DHA and IAA is oxidized to D-dehydroisoascorbic acid (DHIA), cf. structures in Scheme 1. The stoichiometric reactions in the present reaction systems can be expressed by (2a)2IrCl62−+AA⟶2IrCl63−+DHA+2H+(2b)2IrCl62−+IAA⟶2IrCl63−+DHIA+2H+Figure 5
Determination of the reaction stoichiometries by spectrophotometric titrations for reduction of [IrCl6]2– by AA and IAA. Absorbance at 488 nm versus [AA]tot or [IAA]tot for a series of reaction mixtures in which [Ir(IV)] = 0.40 mM kept constant and [AA]tot or [IAA]tot varied from 0 to 0.50 mM. Reaction medium: [H+] = 0.10 M and 1.0 M ionic strength; reaction time is about 5 min, room temperature.


### 3.4. Acid Dissociation Constants of AA and IAA

Acid dissociation constants of AA were reported to be pKa1 = 3.96 [32] and pKa2 = 11.24 [27] at 25.0°C and 1.0 M ionic strength. A value of pKa1 = 2.706 for IAA was reported in solutions of 0.16 M NaCl at 25.0°C [19]; in order to match the reaction conditions used in this work, this value was determined spectrophotometrically in this work at 25.0°C and 1.0 M ionic strength. UV-Vis spectra were recorded for 0.10 mM IAA in a series of buffer solutions covering a pH range from 2.28 to 6.26; before recording each spectrum, a fresh IAA solution and the corresponding buffer solution as reference were thermoequilibrated at 25.0°C for about 10 min. The spectra recorded in pH 2.28, 3.89, and 6.26 buffers are given in the upper panel of Figure 6; an rough isosbestic point in the figure indicates two absorbing species are present. The wavelength of 266 nm was then chosen for the measurements of absorption values. The measured values as a function of pH are given in the lower panel of Figure 6 (data points). Equation (3) was utilized to simulate the data by use of a nonlinear least-squares method [29, 30], where ε1 and ε2 pertain to the molar absorptivities of fully protonated and monoanionic forms of IAA, respectively.(3)Abs266nm=IAAtotε2+ε1⋅10pKa1−pH1+10pKa1−pH.Figure 6
(a) Electronic spectra recorded for 0.10 mM IAA in three different buffer solutions of 1.0 M ionic strength at 25.0°C. (b) Absorption values at 266 nm of 0.10 mM IAA solutions as a function of pH at 25.0°C (data points). The solid line was obtained from the best fit of equation (3) to experimental data by use of a nonlinear least-squares method.The simulation confers a good fit (cf. the lower panel of Figure 6) and affords pKa1 = 3.74 ± 0.05 at 25.0°C and 1.0 M ionic strength; this value is virtually identical to that reported earlier albeit with a difference in ionic strength [19]. Since the pKa2 value was expected to be very close to that of AA and far from the pH region used in this work, we did not determine it.


### 3.5. Reaction Mechanism

The observed second-order rate constants in Table1 increased more than 100 times when pH was changed from 0.16 to 2.33 for both AA and IAA, indicating that their monoanionic forms have a much higher reactivity than the fully protonated forms. The increasing trends are anticipated to continue, but the reactions became too fast to follow by the stopped-flow spectrometer when pH > 2.33. In analogy, the fully deprotonated species of AA and IAA are anticipated to have the highest reactivity among their 3 protolysis species [17]. Since the attributes of the rapid scan spectra suggest that the electron transfer reactions between [IrCl6]2– and AA/IAA undergo directly, it is logical to assume that all the protolysis species of AA/IAA in Scheme 2 will reduce [IrCl6]2–. A reaction mechanism is proposed as delineated in Scheme 2 which involves the three protolysis species of AA/IAA reacting with [IrCl6]2– in parallel; the parallel reactions as indicated by k1-k3 are the rate-determining steps. In each of the rate-determining steps, a single electron transfer takes place, generating a free radical species. Thus, three different free radicals as highly reactive transients are likely involved in the mechanism, cf. the possible structures of the free radicals in Scheme 2 [9, 10, 33]. Each of the free radicals is reacting with another [IrCl6]2– in subsequently fast reaction leading to formation of DHA or DHIA.Scheme 2
Schematically deciphering the reaction mechanism for reduction of [IrCl6]2– by L-ascorbic acid and D-isoascorbic acid (R represents the side chains of AA and IAA, cf. Scheme 1).


### 3.6. Evaluation of the Rate Constants

The rate law of equation (4) can be derived from the reaction mechanism in Scheme 2 in the case of AA, where aH denotes the proton activity corresponding to the pH measurements and the factor 2 in the equation comes from the stoichiometric ratio, cf. equations (2a) and (2b). For the reaction of IAA, the rate law is identical except that [IAA]tot replaces [IAA]tot in (4)−dIrCl62−dt=2k1aH2+k2Ka1aH+k3Ka1Ka2aH2+Ka1aH+Ka1Ka2AAtotIrCl62−.When equation (4) is compared with equation (1a), it confers the expression of k′:(5)k′=2k1aH2+k2Ka1aH+k3Ka1Ka2aH2+Ka1aH+Ka1Ka2.Equation (5) was used to simulate the k′-pH dependence data by use of a weighted nonlinear least-squares method [34]. In the simulation of the AA reaction, pKa1 = 3.96 and pKa2 = 11.24 were used as direct inputs and k1, k2, and k3 were treated as tunable parameters. The simulation resulted in a set of values (k1 = 611 ± 67 M−1s−1, k2 = 2.2 × 107 M−1s−1, and k3 = 3.3 × 1015 M−1s−1); however, the value of k3 is over the diffusion-controlled rate constant (ca. 1010 M−1s−1 in aqueous solution) [35]. Since the k2 >> k1, the k3 value is expected to be also several orders of magnitude higher than k2 [17]. It is thus reasonable to assume that k3 takes a diffusion-controlled rate constant [35]. Given an assumption that k3 = 1010 M−1s−1, the simulation was executed again and the simulated resulted is shown in the upper panel of Figure 4, concurrently providing k1 = 451 ± 59 M−1s−1 and k2 = (2.28 ± 005) x 107 M−1s−1. Alternatively, if the k3-term in equation (5) is neglected, equation (5) is simplified to (6)k′=2k1aH+k2Ka1aH+Ka1.When equation (6) was utilized to simulate the data, an identical fit and identical values of k1 and k2 were obtained (Table 2), indicating that even k3 takes a diffusion-controlled rate constant; the contribution from ascorbate dianion in the pH region used in this work is negligible. This is understandable because the pKa2 value of AA is far from the pH region studied. Nevertheless, a well-defined k2 value is obtained.Table 2
Rate constants for the rate-determining steps at 25.0°C and 1.0 M ionic strength.
ReductantkmValues (M−1s−1)AAk1451 ± 59k2(2.28 ± 0.05) × 107k3Diffusion controlledIAAk1487 ± 45k2(1.19 ± 0.03) × 107k3Diffusion controlledFor the reaction of IAA, the above simulations were also performed and the simulated results were very similar. Consequently, the simulation of thek′-pH dependence data by equation (6) confers an excellent fit (the lower panel of Figure 4) by use of pKa1 = 3.74 determined in this work; the values of k1 and k2 acquired from the fit are listed in Table 2. For both AA and IAA, the relatively large errors associated with k1 values are ascribed to the much higher values of k2, which dominates the total reactivity in the pH region studied.The oxidation kinetics of ascorbic acid by [IrCl6]2– was studied only in a narrow acidity region (0.20 M ≤ [HClO4] ≤ 1.00 M); the derived rate constants were k1 < 200 M−1s−1 and k2 = 1.4 × 107 M−1s−1 at 20.0°C and 1.0 M ionic strength from this region (a correction of a stoichiometric factor was made) [25]. These values were not well determined due to the very narrow acidity region used and are significantly smaller than those obtained in this work; our values derived from a much wider pH region are more accurate and of course more reliable. Moreover, the k3 value is estimated to be a rate constant of the diffusion-controlled one in aqueous solution.


### 3.7. Antioxidative Reactivity of the AA and IAA Species

The acquiredk1 and k2 values for IAA provide a direct comparison with AA. While the k1 values are essentially identical within the experimental errors, the k2 values have a clear difference, namely k2 (AA) ≈ 2k2 (IAA); this reactivity difference is ascribed probably to the slightly lower pKa1 value of IAA. The reactivity difference suggests that the antioxidative power of IAA is weaker than that of AA in the acidic to neutral media where the ascorbate monoanions dominate their respective populations (vide infra).For both AA and IAA, the kinetic characters observed in the present work bolster the conclusion that the redox reactions take place via an outer sphere electron-transfer mode [25, 26]. It is clearly shown from the data in Table 2 that k1 << k2 << k3. In a more visualized way, the species population versus pH distribution diagram and the species reactivity versus pH distribution diagram were contrarily constructed for the AA reaction [36], which are displayed in Figure 7 (the diagrams for the IAA reaction are very similar and are not shown). A few attributes are discerned from the diagrams: (a) the fully protonated AA exists between pH 0 and 6, but it only contributes slightly to the total reactivity in an acidic region (pH from 0 to 2). (b) At the biological matrix of pH 7.0–7.4, the ascorbate monoanion dominates both the species population and the total reactivity. (c) The ascorbate dianion exists only in basic media (pH > 10), but it comes into play as early as pH 6.5 and takes a leading role when pH > 9. The reactivity of AA/IAA species toward reductions of single electron oxidants involved in our body may show similar characters observed above. The significance of the above analysis is that when AA and IAA are used as additives in foods and drinks, these foods and drinks will pass through our digestive systems, where stomach and intestines are in frontlines possessing an acidic environment (pH range from 2 to 5), and the monoanionic forms of AA and IAA may take a leading role in reductions of free radicals or single electron oxidants. Consequently, the antioxidative reactivity of IAA is only about half that of AA.Figure 7
(a) Species of AA versus pH distribution diagram (cf. Scheme 2). (b) Reactivity of AA species versus pH diagram computed by use of k1 = 451 M−1s−1, k2 = 2.28 × 107 M−1s−1 and k3 = 1 × 1010 M−1s−1 (a diffusion-controlled rate constant in aqueous solution). (I) Fully protonated form of AA; (II) monoanionic form of AA; (III) dianionic form of AA.


## 3.1. Spectral Insight into the Reaction Courses

Although an outer-sphere electron transfer was suggested for the oxidation of AA by [IrCl6]2–, the reaction course was not exploited previously [25, 26]. The rapid scan spectra were thus recorded for the oxidations of AA/IAA by [IrCl6]2– in this work; Figure 1 displays such spectra in the case of AA. Clearly the absorption bands around 488, 420, and 306 nm, which are the typical ones of [IrCl6]2–, remained upshifted; in addition, no new absorption bands emerged during the reaction course. Moreover, the kinetic traces recorded at these bands could be well described by single exponentials, giving rise to the values of pseudo-first-order constants kobsd. The above reaction characters indicate that a simple electron transfer took place without complications such as an adduct formation or faster substitution reactions on [IrCl6]2– before the rate-determining step(s). It is thus concluded that the redox reaction is indeed first-order in [Ir(IV)]. The rapid scan spectra in the case of IAA are very similar to those shown in Figure 1.Figure 1
Rapid scan spectra for reduction of [IrCl6]2– by L-ascorbic acid (AA). Reaction conditions: [Ir(IV)] = 0.20 mM, [AA]tot = 2.00 mM, [H+] = 0.200 M, μ = 1.0 M, temperature of 25.0°C. The spectra were acquired at 2, 5, 10, 15, 20, 25, 30, 35, 40, 45, 50, 60, 70, and 80 milliseconds (ms) after mixing.


## 3.2. Second-Order Kinetics

Still under pseudo-first-order conditions, the reaction rates were determined when [AA]tot/[IAA]tot was varied in the region 0.10 ≤ [AA]tot/[IAA]tot ≤ 2.0 mM; for this concentration region, the reaction media had enough buffering capacities to ascertain that no pH changes could be caused by the variations of the concentrations of the reductants. The values of kobsd as functions of [AA]tot or [IAA]tot and of pH were collected; for each concentration, the kobsd value is reported as an average of the at least 5 duplicate runs and the standard deviations were usually less than 5%. Plots of kobsd versus [AA]tot are displayed in Figure 2 and versus [IAA]tot are shown in Figure 3; undoubtedly the linearity of the plots is very good, and no significant intercepts are observable. Hence, the redox reactions are first-order in [AA]tot or [IAA]tot, rendering overall second-order kinetics as described by equations (1a) and (1b), where k′ stands for the observed second-order rate constants.(1a)−dIrIVdt=k′AAtotIrIV,(1b)−dIrIVdt=k′IAAtotIrIV.Figure 2
Dependencies ofkobsd on [AA]tot at 25.0°C and 1.0 M ionic strength for reduction of [IrCl6]2– by AA.Figure 3
Dependencies ofkobsd on [IAA]tot at 25.0°C and 1.0 M ionic strength for reduction of [IrCl6]2– by IAA.Values ofk′ were calculated from the plots in Figures 2 and 3 and are summarized in Table 1. In addition, plots of logk′ versus pH are shown in Figure 4 (data points). The stopped-flow spectrometer was employed essentially to its up limit for following the redox reactions.Table 1
Observed second-order rate constants for the reduction of [IrCl6]2– by AA and IAA at 25.0°C and 1.0 M ionic strength.
ReductantpHk′/M−1s−1AA0.16(7.6 ± 0.1) × 1030.26(9.6 ± 0.2) × 1030.41(1.66 ± 0.02) × 1040.63(2.42 ± 0.04) × 1040.81(3.01 ± 0.06) × 1041.11(5.49 ± 0.09) × 1041.41(1.29 ± 0.03) × 1051.63(2.19 ± 0.05) × 1051.81(3.75 ± 0.09) × 1052.11(6.8 ± 0.2) × 1052.33(1.15 ± 0.06) × 106IAA0.16(7.5 ± 0.1) × 1030.26(8.3 ± 0.2) × 1030.41(1.40 ± 0.02) × 1040.63(2.04 ± 0.04) × 1040.81(2.78 ± 0.06) × 1041.11(4.78 ± 0.08) × 1041.41(9.4 ± 0.03) × 1041.63(1.97 ± 0.04) × 1051.81(3.10 ± 0.07) × 1052.11(6.1 ± 0.1) × 1052.33(1.00 ± 0.02) × 106Figure 4
Second-order rate constantsk′ (in logarithmic scale) versus pH plots at 25.0°C and 1.0 M ionic strength (data points) for reduction of [IrCl6]2– by AA and IAA. The solid lines represent the best fits of equation (6) to the experimental data by use of a weighted nonlinear least-squares method.


## 3.3. The Reaction Stoichiometry

The stoichiometric ratios were determined in a medium of 0.10 M HClO4 and 0.90 M NaClO4 by use of spectrophotometric titration method which was demonstrated to be a very useful approach [28–31]. The absorption values at 488 nm as a function of [AA]tot or [IAA]tot are shown in Figure 5; in either case, the data points all follow two crossing straight lines. The intersection points in Figure 5 confer stoichiometric ratios: Δ[Ir(IV)] : Δ[AA]tot = 0.40 mM : 0.202 mM = 1 : (0.51 ± 0.01) and Δ[Ir(IV)] : Δ[IAA]tot = 0.40 mM : 0.196 mM = 1 : (0.49 ± 0.01). Thus, the derived stoichiometric ratios Δ[AA]tot : Δ[Ir(IV)] and Δ[IAA]tot : Δ[Ir(IV)] = 1 : 2 are established unambiguously when the experimental errors are considered. This is perhaps not surprising because commonly AA is oxidized to DHA and IAA is oxidized to D-dehydroisoascorbic acid (DHIA), cf. structures in Scheme 1. The stoichiometric reactions in the present reaction systems can be expressed by (2a)2IrCl62−+AA⟶2IrCl63−+DHA+2H+(2b)2IrCl62−+IAA⟶2IrCl63−+DHIA+2H+Figure 5
Determination of the reaction stoichiometries by spectrophotometric titrations for reduction of [IrCl6]2– by AA and IAA. Absorbance at 488 nm versus [AA]tot or [IAA]tot for a series of reaction mixtures in which [Ir(IV)] = 0.40 mM kept constant and [AA]tot or [IAA]tot varied from 0 to 0.50 mM. Reaction medium: [H+] = 0.10 M and 1.0 M ionic strength; reaction time is about 5 min, room temperature.


## 3.4. Acid Dissociation Constants of AA and IAA

Acid dissociation constants of AA were reported to be pKa1 = 3.96 [32] and pKa2 = 11.24 [27] at 25.0°C and 1.0 M ionic strength. A value of pKa1 = 2.706 for IAA was reported in solutions of 0.16 M NaCl at 25.0°C [19]; in order to match the reaction conditions used in this work, this value was determined spectrophotometrically in this work at 25.0°C and 1.0 M ionic strength. UV-Vis spectra were recorded for 0.10 mM IAA in a series of buffer solutions covering a pH range from 2.28 to 6.26; before recording each spectrum, a fresh IAA solution and the corresponding buffer solution as reference were thermoequilibrated at 25.0°C for about 10 min. The spectra recorded in pH 2.28, 3.89, and 6.26 buffers are given in the upper panel of Figure 6; an rough isosbestic point in the figure indicates two absorbing species are present. The wavelength of 266 nm was then chosen for the measurements of absorption values. The measured values as a function of pH are given in the lower panel of Figure 6 (data points). Equation (3) was utilized to simulate the data by use of a nonlinear least-squares method [29, 30], where ε1 and ε2 pertain to the molar absorptivities of fully protonated and monoanionic forms of IAA, respectively.(3)Abs266nm=IAAtotε2+ε1⋅10pKa1−pH1+10pKa1−pH.Figure 6
(a) Electronic spectra recorded for 0.10 mM IAA in three different buffer solutions of 1.0 M ionic strength at 25.0°C. (b) Absorption values at 266 nm of 0.10 mM IAA solutions as a function of pH at 25.0°C (data points). The solid line was obtained from the best fit of equation (3) to experimental data by use of a nonlinear least-squares method.The simulation confers a good fit (cf. the lower panel of Figure 6) and affords pKa1 = 3.74 ± 0.05 at 25.0°C and 1.0 M ionic strength; this value is virtually identical to that reported earlier albeit with a difference in ionic strength [19]. Since the pKa2 value was expected to be very close to that of AA and far from the pH region used in this work, we did not determine it.


## 3.5. Reaction Mechanism

The observed second-order rate constants in Table1 increased more than 100 times when pH was changed from 0.16 to 2.33 for both AA and IAA, indicating that their monoanionic forms have a much higher reactivity than the fully protonated forms. The increasing trends are anticipated to continue, but the reactions became too fast to follow by the stopped-flow spectrometer when pH > 2.33. In analogy, the fully deprotonated species of AA and IAA are anticipated to have the highest reactivity among their 3 protolysis species [17]. Since the attributes of the rapid scan spectra suggest that the electron transfer reactions between [IrCl6]2– and AA/IAA undergo directly, it is logical to assume that all the protolysis species of AA/IAA in Scheme 2 will reduce [IrCl6]2–. A reaction mechanism is proposed as delineated in Scheme 2 which involves the three protolysis species of AA/IAA reacting with [IrCl6]2– in parallel; the parallel reactions as indicated by k1-k3 are the rate-determining steps. In each of the rate-determining steps, a single electron transfer takes place, generating a free radical species. Thus, three different free radicals as highly reactive transients are likely involved in the mechanism, cf. the possible structures of the free radicals in Scheme 2 [9, 10, 33]. Each of the free radicals is reacting with another [IrCl6]2– in subsequently fast reaction leading to formation of DHA or DHIA.Scheme 2
Schematically deciphering the reaction mechanism for reduction of [IrCl6]2– by L-ascorbic acid and D-isoascorbic acid (R represents the side chains of AA and IAA, cf. Scheme 1).


## 3.6. Evaluation of the Rate Constants

The rate law of equation (4) can be derived from the reaction mechanism in Scheme 2 in the case of AA, where aH denotes the proton activity corresponding to the pH measurements and the factor 2 in the equation comes from the stoichiometric ratio, cf. equations (2a) and (2b). For the reaction of IAA, the rate law is identical except that [IAA]tot replaces [IAA]tot in (4)−dIrCl62−dt=2k1aH2+k2Ka1aH+k3Ka1Ka2aH2+Ka1aH+Ka1Ka2AAtotIrCl62−.When equation (4) is compared with equation (1a), it confers the expression of k′:(5)k′=2k1aH2+k2Ka1aH+k3Ka1Ka2aH2+Ka1aH+Ka1Ka2.Equation (5) was used to simulate the k′-pH dependence data by use of a weighted nonlinear least-squares method [34]. In the simulation of the AA reaction, pKa1 = 3.96 and pKa2 = 11.24 were used as direct inputs and k1, k2, and k3 were treated as tunable parameters. The simulation resulted in a set of values (k1 = 611 ± 67 M−1s−1, k2 = 2.2 × 107 M−1s−1, and k3 = 3.3 × 1015 M−1s−1); however, the value of k3 is over the diffusion-controlled rate constant (ca. 1010 M−1s−1 in aqueous solution) [35]. Since the k2 >> k1, the k3 value is expected to be also several orders of magnitude higher than k2 [17]. It is thus reasonable to assume that k3 takes a diffusion-controlled rate constant [35]. Given an assumption that k3 = 1010 M−1s−1, the simulation was executed again and the simulated resulted is shown in the upper panel of Figure 4, concurrently providing k1 = 451 ± 59 M−1s−1 and k2 = (2.28 ± 005) x 107 M−1s−1. Alternatively, if the k3-term in equation (5) is neglected, equation (5) is simplified to (6)k′=2k1aH+k2Ka1aH+Ka1.When equation (6) was utilized to simulate the data, an identical fit and identical values of k1 and k2 were obtained (Table 2), indicating that even k3 takes a diffusion-controlled rate constant; the contribution from ascorbate dianion in the pH region used in this work is negligible. This is understandable because the pKa2 value of AA is far from the pH region studied. Nevertheless, a well-defined k2 value is obtained.Table 2
Rate constants for the rate-determining steps at 25.0°C and 1.0 M ionic strength.
ReductantkmValues (M−1s−1)AAk1451 ± 59k2(2.28 ± 0.05) × 107k3Diffusion controlledIAAk1487 ± 45k2(1.19 ± 0.03) × 107k3Diffusion controlledFor the reaction of IAA, the above simulations were also performed and the simulated results were very similar. Consequently, the simulation of thek′-pH dependence data by equation (6) confers an excellent fit (the lower panel of Figure 4) by use of pKa1 = 3.74 determined in this work; the values of k1 and k2 acquired from the fit are listed in Table 2. For both AA and IAA, the relatively large errors associated with k1 values are ascribed to the much higher values of k2, which dominates the total reactivity in the pH region studied.The oxidation kinetics of ascorbic acid by [IrCl6]2– was studied only in a narrow acidity region (0.20 M ≤ [HClO4] ≤ 1.00 M); the derived rate constants were k1 < 200 M−1s−1 and k2 = 1.4 × 107 M−1s−1 at 20.0°C and 1.0 M ionic strength from this region (a correction of a stoichiometric factor was made) [25]. These values were not well determined due to the very narrow acidity region used and are significantly smaller than those obtained in this work; our values derived from a much wider pH region are more accurate and of course more reliable. Moreover, the k3 value is estimated to be a rate constant of the diffusion-controlled one in aqueous solution.


## 3.7. Antioxidative Reactivity of the AA and IAA Species

The acquiredk1 and k2 values for IAA provide a direct comparison with AA. While the k1 values are essentially identical within the experimental errors, the k2 values have a clear difference, namely k2 (AA) ≈ 2k2 (IAA); this reactivity difference is ascribed probably to the slightly lower pKa1 value of IAA. The reactivity difference suggests that the antioxidative power of IAA is weaker than that of AA in the acidic to neutral media where the ascorbate monoanions dominate their respective populations (vide infra).For both AA and IAA, the kinetic characters observed in the present work bolster the conclusion that the redox reactions take place via an outer sphere electron-transfer mode [25, 26]. It is clearly shown from the data in Table 2 that k1 << k2 << k3. In a more visualized way, the species population versus pH distribution diagram and the species reactivity versus pH distribution diagram were contrarily constructed for the AA reaction [36], which are displayed in Figure 7 (the diagrams for the IAA reaction are very similar and are not shown). A few attributes are discerned from the diagrams: (a) the fully protonated AA exists between pH 0 and 6, but it only contributes slightly to the total reactivity in an acidic region (pH from 0 to 2). (b) At the biological matrix of pH 7.0–7.4, the ascorbate monoanion dominates both the species population and the total reactivity. (c) The ascorbate dianion exists only in basic media (pH > 10), but it comes into play as early as pH 6.5 and takes a leading role when pH > 9. The reactivity of AA/IAA species toward reductions of single electron oxidants involved in our body may show similar characters observed above. The significance of the above analysis is that when AA and IAA are used as additives in foods and drinks, these foods and drinks will pass through our digestive systems, where stomach and intestines are in frontlines possessing an acidic environment (pH range from 2 to 5), and the monoanionic forms of AA and IAA may take a leading role in reductions of free radicals or single electron oxidants. Consequently, the antioxidative reactivity of IAA is only about half that of AA.Figure 7
(a) Species of AA versus pH distribution diagram (cf. Scheme 2). (b) Reactivity of AA species versus pH diagram computed by use of k1 = 451 M−1s−1, k2 = 2.28 × 107 M−1s−1 and k3 = 1 × 1010 M−1s−1 (a diffusion-controlled rate constant in aqueous solution). (I) Fully protonated form of AA; (II) monoanionic form of AA; (III) dianionic form of AA.


## 4. Conclusions

The reduction reactions of [IrCl6]2– by AA and IAA have been analyzed kinetically and mechanistically by use of rapid scan and stopped-flow spectral techniques. The reactions strictly follow overall second-order kinetics. The proposed reaction mechanism involves three parallel reactions being the rate-determining steps; in each of the reactions [IrCl6]2– acquires an electron via an outer-sphere electron transfer mode. Rate constants of the rate-determining steps have been derived or estimated which are confidently reliable. A direct comparison between reactivity of AA and IAA demonstrates that IAA has about a half antioxidative reactivity of that of AA in slightly acidic to neutral media. The antioxidative data acquired for AA and IAA can be used as a reference for comparisons with other biologically and biomedically important antioxidants at which we are studying. In addition, pKa1 = 3.74 ± 0.05 at 25.0°C and 1.0 M ionic strength for IAA has been determined in this work.



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*Source: 5505741-2021-11-18.xml*