Copper (II)-Catalyzed Oxidation of Ascorbic Acid: Ionic Strength Effect and Analytical Use in Aqueous Solution

Abstract

Copper is an important metal both in living organisms and in the industrial activity of humans, it is also a distributed water pollutant and a toxic agent capable of inducing acute and chronic health disorders. There are several fluorescent chemosensors for copper (II) determination in solutions; however, they are often difficult to synthesize and solvent-sensitive, requiring a non-aqueous medium. The present paper improves the known analytical technique for copper (II) ions, where the linear dependence between the ascorbic acid oxidation rate constant and copper (II) concentration is used. The limits of detection and quantification of the copper (II) analysis kinetic method are determined to be 82 nM and 275 nM, respectively. In addition, the selectivity of the chosen indicator reaction is shown: Cu²⁺ cations can be quantified in the presence of the 5–20 fold excess of Co²⁺, Ni²⁺, and Zn²⁺ ions. The La³⁺, Ce³⁺, and UO₂²⁺ ions also do not catalyze the ascorbic acid oxidation reaction. The effect of the concentration of the common background electrolytes is studied, the anomalous influence for chloride-containing salts is observed and discussed.

1. Introduction

Copper is one of the most important elements for living organisms, primarily as a catalytic metal in biological systems. It allows many crucial proteins in plants, fungi, and mammalians to function properly [1], including transcriptional regulators, cell transporters and receptors, oxidoreductases, electron transfer proteins, monooxygenases, etc. [2]. Although it is involved in all the biological functions in monovalent form, only Cu²⁺ is bioavailable in significant amounts [3]. The adverse effects of increased copper intake or metal accumulation caused by gene disorder (Wilson’s disease) are also known [1,2]. Copper is a dangerous pollutant of water bodies [4,5] and drinking water, which could be contaminated from copper containers [6]. This shows the importance of copper (II) control in aqueous medium.

Ascorbic acid (AAH₂⁰) is an essential vitamin, which is used for prophylaxis and treatment of scurvy, and it also possesses several other valuable biological properties. Most of these properties are related to the antioxidant properties of AAH₂⁰. For example, its role as a cofactor of many enzymes (e.g., hydroxylases that belong to the Fe²⁺–2-oxoglutarate-dependent families of dioxygenases [7]) is substantiated by the vitamin’s capability of maintaining iron in the ferrous state, thereby ensuring the full activity of this class of enzymes [7]. Another important member of ascorbate-dependent enzymes is the hypoxia-inducible transcription factor (HIF) [7], the upregulation of which is closely related to the development of many forms of cancer [8]. It is not surprising that ascorbic acid was considered a promising compound for preventing or slowing cancer development [9,10,11,12,13], while the oxidized form of AAH₂⁰—dehydroascorbic acid (DHA)—proved useless in inhibiting HIF activity and thus reducing the malignant potential in advanced melanoma [13]. Pro-oxidant action of vitamin C on the complex of copper (II) with 2-phenyl-4-(butylamino)-naphtholquinoline-7,12-dione—which might accelerate the lipid peroxidation in tumors—was also reported [14]. Ascorbic acid shows some antibacterial potential in vitro against both Gram-negative and Gram-positive microorganisms [15] and might be of some use in septic shock treatment; however, rather due to improvement of microvascular function and peripheral tissue perfusion, as well as vasopressor requirement than antibacterial potential (see the recent discussion in Crit. Care [16,17,18,19]). The possible effect of ascorbic acid on wound healing due to proliferative action [20] and the prevention of Alzheimer’s disease due to the efficient binding of AAH₂⁰ to amyloid-β oligopeptide [21] is also being studied.

However, another valuable property of ascorbic acid (also related to the antioxidant activity) is of special interest to us. Ascorbate monoanion (AAH⁻)—which exists predominantly in neutral media considering pK_a1 = 4.14 and pK_a2 = 11.28 [22]—has a low reduction potential (0.19 V for the pair of DHA/AAH⁻ at pH 3.5). However, the rate of redox reaction in absence of catalysts (for example Cu²⁺ [23] or Fe²⁺ ions) is very low as the reaction is spin-forbidden [24]. Only ascorbate dianion (AA²⁻) is capable of auto-oxidation in the air without trace d-metal ion. The fact that the rate of ascorbic acid oxidation depends on the concentration of metal cations is used in analytical chemistry for kinetic determination of some metals concentrations [25,26,27]. Moreover, ascorbic acid consumption is one of the assays used in atmosphere analysis to quantify aerosol oxidative potential, which better predicts adverse health outcomes than particle mass [24,28,29]. This consumption correlates positively with the total concentrations of iron and copper in ambient aerosol [24,30]. Analytical devices that use the pair Cu²⁺/ascorbic acid for the generation of the particles for antioxidant essay [31] and pesticide determination [32], for example, are also known.

An important aspect of both biochemical and analytical applications of ascorbic acid is their complexation with biologically and environmentally relevant cations including Cu²⁺, and such research has been started recently [33,34].

In the present paper, we intend to improve the known technique of quantitative analysis of copper (II) ions using the kinetic method (namely, the kinetics of ascorbic acid oxidation induced by Cu²⁺ ions). In particular: (1) the effect of the ionic strength set by four indifferent electrolytes (NaCl, NaClO₄, KNO₃, KCl) will be tested; (2) the effect of other d-metal ions such as Ni²⁺, Zn²⁺, and Co²⁺ on the possibility of quantitative determination in aqueous solution using ascorbic acid oxidation as an indicator reaction will be studied. It is worth noting that other analytical methods of copper (II) determinations are known, including the classic ones (chelatometry, coulometry, etc.) and several fluorimetric sensors designed recently (see, e.g., just a few of them in papers [35,36,37,38,39,40,41,42,43,44,45,46,47,48]). However, the relatively low sensitivity and selectivity are the disadvantages of the classic techniques, and a preconcentration of copper (II) is required to achieve the nanomolar limit of detection using atomic absorption spectroscopy [49,50,51,52], while the fluorescent sensors are often difficult to synthesize. Moreover, luminescent probes require a non-aqueous solvent for analytical determination, which limits their applicability. Kinetic analysis is as labor intensive as the other methods (since calibration plots are necessary for quantitative analysis in every technique) but also free from the mentioned drawbacks.

2. Results and Discussion

2.1. Quantitative Analysis of Copper (II) Ions Using Ascorbic Acid Oxidation as an Indicator Reaction

The dependence of the conditional rate constant of ascorbate oxidation on the copper (II) concentration is close to linear at any studied ionic strength value, which is set using the various indifferent electrolytes (Figure 1 and Figure S1). ${\mathrm{R}}_{\mathrm{adj}}^2$ varies from 0.92 to 0.99, the worst value of ${\mathrm{R}}_{\mathrm{adj}}^2$ = 0.8624 is observed for the highest concentration of NaCl causing the significant slowing of the oxidation reaction.

The following process is considered for the kinetic measurements:

Ascorbic acid → Dehydroascorbic acid; k′

(1)

where “Ascorbic acid” and “Dehydroascorbic acid” represent all the coexisting protonated species of the reduced (ascorbate) and oxidized (dehydroascorbate) forms of the acid. In other words, C_total(AAH₂⁰) = [AAH₂⁰] + [AAH⁻] + [AA²⁻]; C_total (DHA) = [DHA⁰] + [DHA⁻].

These calibration plots (Figure 1 and Figure S1) allow one to find the unknown concentration of copper (II) ions. The limits of detection (LOD) and quantification (LOQ) could be calculated using the standard 3σ and 10σ rule [53]. However, since the analytical response in this case is a conditional rate constant instead of absorbance or fluorescence intensity, a corresponding modification of the method [53] is required. The lowest rate constant that could be calculated from the UV–vis spectral data corresponds to the lowest detectable change in the absorbance during 600 s of kinetic experiment. In the absorbance range of the 0 to 1, the error of experimental A measurements claimed by the manufacturer is 0.002 units. Therefore, the lowest detectable change in absorbance—which can be referred to the catalytic action of the copper (II) ions—is 0.004 units (because the error can be made at the beginning and the end of the kinetic experiment). The lowest rate constant corresponding to that low alteration of A is equal to 6.68·10⁻⁶ s⁻¹, which gives a LOD of 82 nM and a LOQ of 275 nM at zero ionic strength considering k′_cat = 243.08 s⁻¹ (k′_cat, the rate constant of the catalytic reaction is equal to the rate of hypothetical reaction of ascorbic acid oxidation, where C_cat = 1 mol L⁻¹, and is equal to the slope in Figure S1a). In our opinion, it is a good result in comparison even with more complex fluorimetric techniques using sophisticated chemosensors. Although the kinetic method using ascorbic acid oxidation shows higher LOD for Cu²⁺ than 20 nM [35], 4.7 nM [36], and 24.5 nM [37] and slightly higher than LOD of inductively coupled plasma-atomic emission spectroscopy (74 nM [54]), this technique is more sensitive than the reported fluorescent ones [38,39,40,41,42,43,44,45,46,47,48], atomic absorption spectroscopy (AAS) [49] (however, preconcentration methods [50,51,52] improves the sensitivity of AAS by 10 times compared to the proposed method) and atomic emission spectroscopy [54]). Another advantage of the proposed method is its probable applicability for determining the equilibrium concentration of copper (II) ions in complex systems containing ligands (including biomacromolecules such as nucleic acids and proteins), i.e., for studying the coordination equilibria and determining the stability constants of Cu²⁺ complexes. It should be noted, however, that the interactions between copper (II) ions and biomacromolecules can be studied only in an absence of other ligands, in the limited range of ionic strength value and using the buffer, which does not bind metal ions into complexes.

The LOD and LOQ can be further improved if the small ionic strength value is set using chloride-containing electrolyte, which leads to a steeper slope.

2.2. Effect of the Nature and Concentration of Background Electrolytes on the Rate of Ascorbic Acid Oxidation

The initial set of experiments, with NaCl utilized as an indifferent electrolyte, showed that the effect of salt concentration on the apparent rate constant of vitamin C auto-oxidation is non-linear (Figure 2a). Moreover, the decrease in the NaCl content did not make the observed value of k′ closer to that determined for zero ionic strength. The first small addition of this background electrolyte to the system accelerates the indicator reaction, and only I(NaCl) ≥ 0.1 mol L⁻¹ decreases the rate constant. We performed additional experiments with other often used background electrolytes (KCl, NaClO₄, KNO₃) to figure out whether it is the common property of the electrolytes or the unique peculiarity of NaCl or every salt bearing chloride ion.

Our results give some evidence that the latter is true: 0.05 M KCl makes the indicator reaction of ascorbic acid autoxidation as quick as 0.05 M NaCl (Figure 2b). Other salts used to set the ionic strength value also give the non-linear dependence of k′ on the electrolyte concentration (Figure 2c,d).

The oxidation of ascorbic acid induced by copper (II) ions is a complex sequence of the reactions influenced in a different manner by the ionic strength. The shifts of chemical equilibria caused by the variation in concentration of background electrolyte may also change the equilibrium concentrations of components involved, both directly (since equilibrium constants depend on ionic strength value; primary salt effect) and indirectly (due to the reactions between background electrolyte and the involved components; secondary salt effect). Zhou et al. [23] give the general mechanism of vitamin C oxidation in the presence of Cu²⁺ ions; however, the most detailed description of the reactions occurred in the studied system is provided in a report [24]. Using this complex chemical model, the authors [18] determined the apparent rate constants for the following general processes.

Cu²⁺′ + AAH₂⁰ + O₂ = Cu²⁺′ + DHA + H₂O₂; k′ = 7.7 × 10⁴

(2)

Cu²⁺′ + AAH⁻ + O₂ = Cu²⁺′ + DHA + H₂O₂ − H⁺; k′ = 2.8 × 10⁶

(3)

where DHA is dehydroascorbic acid, Cu²⁺′ represents all species involving copper (II) ions including free Cu²⁺ ions and CuOH⁺, CuCl⁺, and CuCl₂ as well as Cu²⁺ complexes with ascorbic acid of different composition [34,55,56] etc.

The secondary salt effect leading to the change in ascorbic acid dissociation degree can be neglected for two reasons. First, the variation in the dissociation constants of vitamin C with the ionic strength alteration up to 0.25 mol L⁻¹ (NaCl) is negligible [22]. AAH₂⁰ and AAH⁻ coexist in roughly equimolar quantities when the ascorbic acid is dissolved in the distilled water [22]. Second, the monoascorbate anion undergoes oxidation much more readily than ascorbic acid [24]. Therefore, the loss in its concentration would be refilled via instantaneous dissociation of AAH₂⁰ caused by the shifted chemical equilibrium between two species. There are limited data on dehydroascorbic acid dissociation [57]; however, it is unlikely that a small change in ionic strength value could significantly alter the pH value of the solution.

The secondary salt effect on the composition of the mixture of different copper (II) species is more complex.

First, the variation of the ionic strength may influence the coordination equilibria between Cu²⁺ and ascorbic acid considering a relatively large yield of the complex under our experimental conditions (1.6 × 10⁻⁶ to 3.4 × 10⁻⁵ mol L⁻¹ of [CuAAH₂]²⁺ for C⁰(Cu²⁺) = 4.3 × 10⁻⁶ and 1.06 × 10⁻⁴ mol L⁻¹, respectively, taking into account the stability constants from [41,42]). If the results of Ritacca et al. are used for the calculation of equilibrium compositions, the predominant complexes are [CuAA₃]⁴⁻ when copper (II) ions is taken in a short supply, and [Cu(OH)AA]⁻ if the total concentration of Cu²⁺ exceeds that of ascorbic acid. The catalytic action of the complexes is probably less pronounced than that of free copper (II), and alteration of the ionic strength leading to the shift of coordination equilibria towards the reactant should accelerate the reaction of auto-oxidation of vitamin C, while the stabilization of the complexes should slow the process of AAH₂⁰ destruction.

Second, NaCl is truly indifferent only towards Cu²⁺ ions as the chloride complexes of copper (II) are very weak [24,58]. It does not apply to the chloride complexes of copper (I) [59], which plays an important role in the Fenton-like process according to the following equations [24]:

Cu²⁺′ + HO₂· = Cu⁺ + O₂ + H⁺; k′ = 10⁸

(4)

Cu²⁺′ + H₂O₂ = Cu⁺ + O₂⁻ + 2H⁺; k′ < 1 for Cu²⁺, CuOH⁺ and CuSO₄;
k′ = 70 for CuCl⁺ and CuCl₂

(5)

Cu⁺ + O₂ ⇋ Cu²⁺ + O₂·−; k_1′ = 4.6·10⁵; k_{−1′ =} 8 × 10⁹

(6)

Cu⁺ + OH· = Cu²⁺ + OH⁻; k′ = 3 × 10⁹

(7)

Cu⁺ + H₂O₂ = Cu²⁺ + OH· + OH⁻; k′ < 100

(8)

Cu⁺ + H₂O₂ = Cu³⁺ + 2OH⁻; k′ = 61

(9)

Cu⁺ + Cu³⁺ = 2Cu²⁺; k′ = 3.5 × 10⁹

(10)

Cu⁺ + HO₂· = Cu²⁺ + H₂O₂ − H⁺; k′ = 2.3 × 10⁹

(11)

Cu⁺ + O₂·⁻ = Cu²⁺ + H₂O₂ − 2H⁺; k′ = 9.4 × 10⁹

(12)

The rate constants use s⁻¹ time scale.

Chloride complexes of copper (I) also undergo the oxidation [60,61]:

CuCl + H₂O₂ = CuCl⁺ + OH⁻ + OH·; k′ = 3.7 × 10³

(13)

CuCl₂⁻ + H₂O₂ = CuCl⁺ + OH⁻ + OH· + Cl⁻; k′ = 190

(14)

CuCl + O₂ = Cu²⁺ + O₂⁻; k′ = 19

(15)

CuCl₂⁻ + O₂ = Cu²⁺ + O₂⁻; k′ = 0.12

(16)

CuCl₃²⁻ + O₂ = Cu²⁺ + O₂⁻; k′ = 0.44

(17)

Gonzales-Davila et al. [62] deduced an equation expressing the correlation between the logarithm of the rate constant of copper (I) oxidation and pH, temperature, and ionic strength value. This equation appears to be valid only within the parameters used for its derivation (pH 7.2 to 8.5; T of 278.2 to 308.2 K; I(NaCl) 0.1 to 0.7 mol L⁻¹). However, under different experimental conditions, at constant pH and temperature it may also have the form proposed [62]:

$$\lg {\mathrm{k}}^{\prime }=\lg {{\mathrm{k}}_0}^{\prime }+B\sqrt{I}+CI$$

(18)

where k′ and k_0′ are the rate constants at given and zero ionic strength, respectively; I is an ionic strength value, B and C are the empirical constants.

We tried to use Equation (18) to describe the experimental data (Figure 2a), and the expression also seems valid to the results (Figure 3):

This might be an indirect indication of the fact that the anomalous chloride concentration effect on the conditional rate constant of vitamin C oxidation in the presence of copper (II) ions might be explained by the changes in rate constant of copper (I) oxidation. However, no less than three objections could be raised against this conclusion: (i) taking the logarithm of the dimensional value is, strictly speaking, prohibited; (ii) taking the logarithm is recognized as some sort of mathematical trickery (see, e.g., [63]), which helps greatly in smoothing the dependence (which, otherwise, might be indicative of additional factors influencing it); (iii) the possibility of the same equation use for two dependencies does not mean any link between them (although does not exclude such an option either).

Other indifferent electrolytes (NaClO₄, KNO₃) cause a decrease in the apparent rate constant of ascorbate autooxidation induced by copper (II) ions. It can indicate that the rate-limiting stage of this reaction is an interaction between ions with the opposite charges, and the primary salt effect slows the overall process.

2.3. Selectivity of Copper (II) Quantitative Analysis Using Ascorbic Acid Oxidation as an Indicator Reaction

We tested the effect of cations such as Ni²⁺, Co²⁺, Zn²⁺, La³⁺, Ce³⁺, and UO₂²⁺ on the apparent rate constant of ascorbic acid, and observed none. These cations do not catalyze the destruction of vitamin C. Figure 4 shows the possibility of using ascorbic acid for the copper (II) concentration determination when some other cations are also present in the solution:

The differences between dependencies of the conditional rate constant on the copper (II) concentration with and without additions of other d-metal ions are negligible. They do not interfere with the quantitative analysis of Cu²⁺ in an aqueous solution.

Finally, we tried to analyze the aqueous solution containing simultaneously Ni²⁺, Co²⁺, Zn²⁺ (concentration of each 10⁻⁴ mol L⁻¹), and Cu²⁺ of 2.15 × 10⁻⁵ mol L⁻¹ concentration. Three experiments returned the value of copper (II) concentration 2.12 ± 0.11 × 10⁻⁵ mol L⁻¹, which corresponds to a recovery of 98.6 ± 5.1%.

3. Materials and Methods

3.1. Chemicals

Ascorbic acid (Khimreaktiv, Nizhnii Novgorod, Russia), Cu(NO₃)₂ (Acron Organics, New Jersey, USA), NaCl, KCl, NaClO₄, KNO₃, Cu(NO₃)₂, Ni(NO₃)₂, Co(NO₃)₂, Zn(NO₃)₂ (all Khimreaktiv, Nizhnii Novgorod, Russia), La(NO₃)₃, Ce(NO₃)₃, and UO₂(CH₃COO)₂ (all RedkiiMetallRF, Novosibirsk, Russia) were used without additional purification. Copper (II), nickel (II), cobalt (II), zinc (II), lanthanum (III), cerium (III), and uranyl (VI) concentrations in the stock 0.01 M solutions were determined chelatometrically. Only freshly prepared solutions of ascorbic acid were used. The concentration of ascorbic acid was determined spectrophotometrically using the literature values of the molar extinction coefficients [22]. All solutions were prepared using bidistilled water (κ = 3.6 μS/cm, pH = 6.7). The necessary amounts of background electrolytes were weighted using Shimadzu laboratory balances (AUW220, Shimadzu, Japan) with an error of weighing of 5 × 10⁻⁵ g.

3.2. Copper (II) Assay

An aliquot (1 to 25 μL) of aqueous solution of Cu(NO₃)₂ of the concentration of 0.01161 mol L⁻¹ concentration, which corresponds to the final concentration of 4.3 × 10⁻⁶ to 1.06 × 10⁻⁴ mol L⁻¹ was added to 2.7 mL of aqueous solution of ascorbic acid (1.5 × 10⁻⁴ mol L⁻¹) and the necessary background electrolyte placed in a standard quartz cell with an optical path length of 1 cm, and quickly shaken. The absorbance of the resulting mixture was then recorded every 10 s during 600 s at a wavelength of 265 nm, which corresponds to the maximum absorbance of ascorbic acid and negligible absorbance of dehydroascorbic acid and other possible products/semi-products. Kinetic measurements were performed on the Shimadzu1800 double-beamed spectrophotometer (Shimadzu, Columbia, MD, USA).

The dependence of absorbance on time was then recalculated using the literature values of molar extinction coefficients of ascorbic acid [22], and the resulting data ‘ascorbic acid concentration’ vs. ‘time’ were processed using Kinet software [64] to obtain the apparent rate constant of the ascorbic acid oxidation. All kinetic experiments were at least triplicated; the error bars in Figure 1, Figure 2, and Figure 4 represent the standard deviations calculated for the corresponding sample size.

The rate of the catalytic monomolecular reaction is expressed as

$$\frac{d{C}_A}{d\tau }=-{\mathrm{k}}^{\prime }{C}_A$$

(19)

where

k′ = k₀ + k_cat C_cat

(20)

where k′ is an observed apparent rate constant; k₀ is a rate constant of the reaction in the absence of catalyst; k_cat is a rate constant of the catalytic reaction; C_cat is a catalyst concentration.

As it follows from Equation (20), the observed rate constant is linearly dependent on the catalyst concentration.

4. Conclusions

In the present contribution, we gave a further development of the kinetic analytical technique for quantitative determination of copper (II) ions using copper-induced autooxidation of ascorbic acid as an indicator reaction. The method is cheaper in comparison with copper (II) analysis using the fluorescent chemosensors. Moreover, it does not require non-aqueous solvents. In addition, kinetic analysis shows limit of detection and quantification as low as those characteristic for some fluorimetric techniques (82 and 275 nM, respectively). Kinetic analysis is also selective, allowing Cu²⁺ detection in solutions containing Ni²⁺, Co²⁺, and Zn²⁺ ions as well.

The effect of the ionic strength set by four different background electrolytes (NaCl, KCl, KNO₃, NaClO₄) on the rate of oxidation of vitamin C is also studied. Chloride-containing salts taken at low concentrations (up to 0.1 mol L⁻¹) accelerate the ascorbate destruction reaction, while a higher electrolyte content causes slowing of the process. This can probably be explained by the formation of chloride complexes of copper (I), whose oxidation rate differs from that of the free cation.

We believe that the perspective of the present work is using the kinetic analysis to study the interactions of copper (II) ions with biomacromolecules (proteins, DNA), which is difficult to perform using other methods. However, the limits of the proposed technique should be kept in mind: (1) other ligands capable of Cu²⁺ binding should be absent; (2) buffer agent should be indifferent towards cation; (3) the possible ionic strength value is limited (which is not the biggest problem because the blood serum ionic strength is 0.15 mol L⁻¹).