Next Article in Journal
Teachers Perceptions on Early Childhood’s Traffic and Life Safety Education Program Using VR
Next Article in Special Issue
The Effect of Oxygen Tension on the Differentiation of Outgrowth Cells from Embryoid Bodies Produced by Mouse Induced Pluripotent Stem Cells
Previous Article in Journal
Revalidation Technique on Landslide Susceptibility Modelling: An Approach to Local Level Disaster Risk Management in Kuala Lumpur, Malaysia
Previous Article in Special Issue
Machine Learning in Bioequivalence: Towards Identifying an Appropriate Measure of Absorption Rate
 
 
Search for Articles:
Title / Keyword
Author / Affiliation / Email
Journal
Article Type
 
 
Advanced Search
 
Section
Special Issue
Volume
Issue
Number
Page
 
Journals
Applied Sciences
Volume 13
Issue 2
10.3390/app13020776
applsci-logo
Submit to this Journal Review for this Journal Propose a Special Issue
► Article Menu

Article Menu

share Share announcement Help format_quote Cite question_answer Discuss in SciProfiles thumb_up
...
Endorse textsms
...
Comment
first_page
settings
Order Article Reprints
Font Type:
Arial Georgia Verdana
Font Size:
Aa Aa Aa
Line Spacing:
Column Width:
Background:
Article

Wide pH Range Potentiometric and Spectrophotometric Investigation into the Acidic Constants of Quercetin, Luteolin and l-Ascorbic Acid in Aqueous Media

by
Luana Malacaria
and
Emilia Furia
*
Department of Chemistry and Chemical Technologies, University of Calabria, 87036 Rende, Italy
*
Author to whom correspondence should be addressed.
Appl. Sci. 2023, 13(2), 776; https://doi.org/10.3390/app13020776
Submission received: 25 November 2022 / Revised: 26 December 2022 / Accepted: 1 January 2023 / Published: 5 January 2023
(This article belongs to the Special Issue New Trends in Biosciences III)
Browse Figures
Review Reports Versions Notes

Abstract

:
It is now well established that the dissociation constants of an organic compound are characteristic of the types of groups, or the combinations of groups, contained in it. Furthermore, the acid–base dissociation constants are important parameters to fully understand the properties of a molecule in biological systems. In this framework, the aim of the present study was to determine the acidic constants of three natural molecules with well-known antioxidant properties, namely quercetin, luteolin and l-ascorbic acid. The evaluation was carried out in aqueous media (i.e., 0.16 M NaCl) at 37 °C in a wide pH range by using a combined approach based on potentiometric and spectrophotometric measurements. The results underline the necessity to employ both experimental techniques to obtain accurate values for acidic constants, preventing uncertainties related to undesirable oxidation reactions.

1. Introduction

The thermodynamic properties of bioactive molecules in aqueous solutions are of significant scientific and practical interest. In this context, the acidic constants are important physicochemical parameters of a molecule, as they reflect its reactivity, solubility, affinity for a macromolecule and/or ability to establish intermolecular interactions. In addition, the acid and basic functional groups of a molecule strongly influence its pharmacokinetics and toxicity. The experimental evaluation of the pKa values of the molecules of biological interest, which could be hygroscopic and very unstable in an extreme pH range, is still unsolved, as it is related to the conditions in which the experiments are performed. The most useful and widely employed conventional methods are potentiometric and spectrophotometric titrations and conductometry [1,2,3]. However, only a few of them have referred to measurements conducted in aqueous media, where they show low water solubility.
In this framework, the present study intends to evaluate the acidic constant values of quercetin, luteolin and l-ascorbic acid (Figure 1a–c), three natural antioxidants (generically HnL) whose beneficial properties are well known, in aqueous solution. The experimental approach provides measurements via potentiometric and UV-vis spectrophotometric titrations in a wide pH range, at 37 °C and ionic strengths of 0.16 M, using NaCl as a background electrolyte. The constant ionic medium method has proved to be indispensable in equilibrium studies of complex ionic reactions as it is necessary to minimize activity coefficient variation, even though the modification is in the concentration of the reagent. The method, which consists of using concentrated aqueous solutions of inert salts as the solvent, allows for the use of concentrations instead of activities in the evaluation of equilibrium constants. Additionally, it was useful to minimize the liquid junction potential, owing to the hydrogen ion concentration, which was varied in a wide pH range to assess the acidic constants of HnL.
Quercetin and luteolin belong to flavonoids that are secondary plant metabolites of the phenolic family. The most common flavonoids contain three rings, indicated as A, B and C (Figure 1a,b), and the oxidation state of C ring provides their classification into several subgroups, as quercetin and luteolin belong to flavonol and flavone subclasses, respectively. l-ascorbic acid is a water-soluble vitamin required for the prevention of scurvy and is known to be active as a cellular hydrogen transfer carrier for the redox enzyme in live cells [4]. Structurally, ascorbic acid is a γ-lactone and an ene-diol (Figure 1c).
All these plant-derived natural compounds are found in significantly high quantities in fruits and vegetables and are considered valuable health-promoting compounds [4,5,6,7,8,9]. Choosing these natural antioxidants is also related to their well-known capability to chelate several metal ions of relevant importance for biological systems, and they have attracted our research interest in the past 10 years [10,11,12,13,14,15,16,17]. Their ability to bind cations depends on their structure, characterized by functional groups such as phenol, enol and enediol moieties that can undergo deprotonation via metal interaction, whatever their pKa value. In this context, the accurate knowledge of the acidity of -OH is fundamental to understanding how these molecules can counteract the metal-related damage [18,19,20,21,22,23,24,25]. Furthermore, their antioxidant ability could be correlated to the relative acidity of -OH groups.

2. Materials and Methods

2.1. Chemicals

All chemicals were of analytical grade. The titrant carbonate-free solutions of sodium hydroxide, the hydrochloric acid and the sodium chloride stock solutions were prepared and standardized as previously described [12].
Quercetin (Sigma, ≥95%), luteolin (Sigma, ≥98%) and l-ascorbic acid (Sigma, 99.2%), were stored in a desiccator over silica gel and used as such.
To avoid variations during dilution, all the solutions were prepared by adding the appropriate amount of NaCl, as the background electrolyte, to set the ionic strength at 0.16 M. Double-distilled water, freed from any organic impurities by means of a Milli-Q system (Millipore), was used to prepare fresh solutions.

2.2. Potentiometry and Spectrophotometry

The potentiometric arrangement and the titrations were conducted with the same apparatus described previously [12]. The glass electrodes were from Metrohm. Following the addition of the reagents, they acquired a constant potential (within ±0.01 mV) by 30 min. To prevent any interference due to carbonate formation, nitrogen gas was gently introduced into the test solutions, which were magnetically stirred throughout titrations. Nitrogen commercial cylinders were used, and the gas was further purified by passing it through 1 M NaOH, 1 M H2SO4, double-distilled water and 0.16 M NaCl. The cell apparatus was kept in a thermostatic bath at (37.0 ± 0.1) °C.
The acid–base equilibria were investigated by performing potentiometric titrations, at 37 °C and in 0.16 M NaCl, and by measuring, with a glass electrode (GE), the electromotive force (EMF) of cell:
REǀTest SolutionǀGE
where RE corresponds to the reference electrode (Ag|AgCl|0.01 M NaCl|0.16 M NaCl), while the test solution was made of CL M of HnL, CA M HCl, CB M NaOH and 0.16 M NaCl. The primary CL, CA, CB and [H+] represent the basic data in the treatment to obtain the acidic constants.
The EMF of the cell can be expressed, in mV, at the temperature of 37 °C, as Equation (1):
E = E° + 61.54 log[H+] + Ej
where E° was constant for each series of measurements and Ej is the liquid junction potential, which depends on [H+] only [26]. The Ej value at 37 °C and 0.16 M NaCl was determined by acid–base titration using cell (G) without HnL. The obtained results for [H+] ≤ 0.100 M could be accurately fitted by the linear slope Ej ([H+])I = −j I [H+]. The value of j (mV/M) was 631 ± 1.
All titrations were separated into two parts. In the first one, E° constant value (within ±0.05 mV), was determined in accordance with the Gran’s method [27,28] in the [H+] range 10−4–10−2 M and without HnL. [H+] was stepwise reduced by coulometrically generating OH− ions with the circuit C:
− Pt/Test Solution/AE +
where AE is the auxiliary electrode of generical composition: 0.16 M NaCl/0.1 M NaCl, (0.16–0.1) M NaCl/Hg2Cl2/Hg.
Under the chosen experimental condition, the only electroreducible species was water. Hence, suppose that at the cathode the only reaction that occurs is the water reduction (Equation (2)):
H2O + e− ⇄ ½ H2 + OH−
In the test solution of a known volume, V, CB can be expressed as (µF·10−6/V) M, where µF stands for the microfaradays passed through the cell.
In the second part, after the addition of HnL, the acidity was gradually decreased by adding known volumes of the standard NaOH solution.
All titrations were conducted with a programmable computer-controlled data acquisition switch unit 34,970 A from Hewlett and Packard. The constant–current source of the coulometric circuit was a system DC power supply 6614 C by Hewlett and Packard. The EMF values were measured at a precision of ±10−5 V using an OPA 111 low-noise precision DIFET operational amplifier.
Given the low solubility of quercetin and luteolin in water [12,14] and by using the same experimental approach tested in a previous study [29], all the titrations were carried out by adding an exactly known and weighed amount of solid in the cell apparatus. When equilibria (i.e., Equations (2) and (3)) take place, quercetin and luteolin dissolved into the aqueous medium. CL varied between 0.5 and 5 × 10−3 M, while the hydrogen ion concentration varied from 5.0 × 10−3 M (pH 2.3) to 1.6 × 10−9 M (pH 8.8). The higher pH limit was imposed on the accuracy in the EMF measurements by glass electrode. For this reason, all systems were studied by UV-vis spectrophotometric titrations to investigate the more alkaline pH range (i.e., up to pH 12). CL was varied between 0.01 and 0.04 × 10−3 M. The spectrophotometric measurements were carried out with a Varian Cary 50 Scan UV Visible Spectrophotometer, by controlling the temperature of the cell holder at (37.0 ± 0.3) °C by using a Grant circulating water bath. Matched quartz cells of thickness 1 cm were employed.
The absorbance values, Aλ, recorded to 0.001 units between 200 and 600 nm for quercetin and luteolin, and between 200 and 340 nm for l-ascorbic acid, were measured each 2 nm to collect data for numerical elaboration by using a Hyperquad [30]. The spectra were recorded 2 minutes after mixing with a strong base. This period was thought short enough to avoid any photochemical degradation. The acquisition of data was controlled with the aid of a computer linked to the instrument.

3. Results

Potentiometry and spectrophotometry were used to determine the acidic constants of quercetin, luteolin and l-ascorbic acid—generically HnL. To acquire accurate results, the physical integrity of the molecules and the stability at different pH values in which the study was carried out must be taken into consideration during the whole experiment. So the systems were kept under an inert atmosphere and prevented from being exposed to light radiation.
The general equilibria reported in Equations (3) and (4) fully explain the potentiometric data, treated by numerical and graphical methods:
HnL + H3O+ ⇄ Hn+1L+ + H2O    log K* (±3σ)
HnL + H2O ⇄ Hn−mLm− + m H+    log Kam (±3σ)
These correspond to the protonation of the carbonyl moiety, according to [31], and to the protolysis of the -OH groups, respectively. The data used to determine the acidic constants, according to Equations (3) and (4), were obtained by conducting two titrations for each compound.
The interpretation of the potentiometric experimental points, in terms of species and equilibrium constants, was carried out with the least-squares computer program Superquad [32] to get the minimum of the function reported by Equation (5):
U = ∑(Eiobs − Eical)2
The results are reported in Table 1.
The graphical evaluation was carried out by analyzing the trend of experimental points, depicted as ZH and obtained at different CL values, with respect to the pH. ZH represents the mean number of protons released for ligand [33], and it is equal to ([H+]–CA + CB + Kw/[H+])/CL. The values for the ion product of water, in our experimental condition (i.e., 37 °C and 0.16 M NaCl), was taken from [34].
The results are reported in Figure A1a–c, for quercetin, luteolin and l-ascorbic acid, respectively. For different CL values, the experimental points overlap, within the limit of experimental error.
As can be seen in Figure A1, the data are fully explained by Equations (3) and (4). In particular, at pH values lower than 3, ZH tends to be −1 for all the investigated systems, confirming that a protonation occurs, according to Equations (3) and (4).
Furthermore, as expected, the protolysis equilibria of quercetin and luteolin occur in the alkaline range (−log [H+] > 8) at the expenses of the hydroxylic groups (see Figure 1a,b). In contrast, for l-ascorbic acid, the dissociation starts at a pH higher than 3 (Figure A1c) with the protolysis of -OH in position 3 (see Figure 1c).
Accurate values for the constants of protolysis equilibria (Equation (4)) were determined for quercetin and luteolin only. By performing a back titration, whose points are depicted in Figure A1c as orange crosses, we have verified that l-ascorbic acid undergoes a degradation in an alkaline medium (i.e., pH higher than 7.5), which occurs faster than the potentiometric measurement time.
Therefore, we have performed UV-vis spectrophotometric titrations at different CL values (i.e., between 0.01 and 0.04 mM). UV-vis absorption spectrophotometry is a simple and accurate method for determining acidic constants, whose determination is based on a change in the shape and intensity of the absorption spectra with a change in the pH of the test solutions.
The region of pH close to the pKa values of corresponding acids was supposed to be optimal to obtain accurate values, in the choice of the pH range for all the compounds. Additionally, the spectral range for all compounds was designated according to the variations in the spectra at both acidic and basic solutions.
Consequently, the hydrogen ion concentration was varied by adding small and exactly known volumes of a standard NaOH solution (CB), from 1.0 × 10−5 M to 1.0 × 10−12 M for quercetin and luteolin and from 1.0 × 10−2 M to 1.0 × 10−11 M for l-ascorbic acid, and the absorbance values were recorded at 93 wavelengths between 230 and 457 nm for quercetin and luteolin and at 50 wavelengths between 211 and 330 nm for l-ascorbic acid.
The typical recorded spectra are depicted in Figure A2a–c, for quercetin, luteolin and l-ascorbic acid, respectively.
The UV-vis spectra of quercetin and luteolin (Figure A2a,b, respectively) show two main absorption bands, which are comparable to signals observed for most flavonoids [12,14]: band I in the 350–400 nm range, as a result of the conjugation between the B and C ring (cinnamoyl system), and band II in the 240–270 nm range, deriving from the A-C ring (benzoyl system). Moreover, two other less-intense absorption bands at 303 and 280 nm for quercetin and at 295 and at 267 nm for luteolin were observed. The UV-vis spectrum of l-ascorbic acid exhibits a single band centered at 260 nm, in accordance with the literature [11,15].
Figure A2a shows that the maximum of band I was blue shifted (i.e., from 367 to 320 nm) by moving from pH 5.5 to pH 6.5. The absorption profile changed negligibly upon raising the pH from 6.5 to 12.0, indicating, according to [35], that the acid–base sites of quercetin were not appreciably rearranged during the dissociation in this pH range.
In contrast, for luteolin and l-ascorbic acid (Figure A2b,c) the maxima of the principal absorption bands were red shifted (i.e., from 352 to 402 nm and from 243 to 267 nm, for luteolin and l-ascorbic acid, respectively) by moving from acidic to basic pH, attesting to the formation of a new spectrally active species in the solution.
The numerical elaboration of spectrophotometric data was carried out by simulating each titration from the knowledge of the total concentrations, CB and CL. The minimum of the function, described by Equation (6), was sought by employing the Hyperquad program [30]:
U = ∑i∑k wk (Aik − Aikc)2
The ion product of water was taken from [34] and was kept invariant.
The results are given in Table 1.

4. Discussion

It has been suggested in previous studies that the antioxidant ability of a series of compounds is strongly affected by pH, which influences the relative abundance of the corresponding acid–base forms [36]. Acid–base equilibria have been reported to strongly affect the UV-vis spectra of several phenolics, including flavonoids. This is observed as significant changes in both the shapes and the intensities of the bands owing to variation in the pH values [37]. The absorption and fluorescence spectra of luteolin have been reported to strongly depend on pH [38]. Therefore, the corresponding acidity constant values are an essential parameter in predicting the functional mechanism of proteins in chemistry and biochemistry.
We calculated here the acidic constants of all the possible sites of quercetin, luteolin and l-ascorbic acid. According to our results, the increase in the pH of the medium at values higher than 3 prompted the proton removal from the protonated carbonyl group (i.e., from Hn+1L+ in Equation (3)), leading to the formation of an electroneutral compound (generically HnL). In contrast, in the narrow pH range of 8–10, in an aqueous solution of quercetin and luteolin, six or five species can simultaneously exist, respectively. These ion molecular forms are in dynamic equilibrium in the aqueous solution, depending on the acidity of the medium. The experimental assignment for the constants of quercetin and luteolin could not be assessed. Although some efforts have been carried out in this way, a systematic and detailed study factoring in all the possible acid forms has not been carried out, and the computational evaluation is still debated [31,35,36,37,38,39]. On the basis only of their structures and, in particular, on the number of phenolic -OH groups, the existence of five and four pKa values for quercetin and luteolin, respectively, can be hypothesized. However, for quercetin, in previous works dealing with its acidity constants, sets of only two [40,41], three [41,42,43,44] and four [43] pKa values have been reported. Interestingly, in the work of Álvarez-Diduk and colleagues [36], through a combined experimental and computational approach, a complete set of constants has been reported for quercetin. In particular, the deprotonation order was supposed to be 4′, 7, 3, 3′ and 5. Luteolin is structurally analogous to quercetin—except that it has four phenolic -OH groups, as there is no -OH group on ring C. Indeed, the Ka1 and Ka2 of quercetin and those of luteolin are the same. Hence, we can suppose that the first two protolysis equilibria involve -OH groups that are less influenced by the deprotonation on the C ring. Only the first acidity constant has been experimentally measured in an aqueous medium [14], and the -OH group in position 4′ has been theoretically calculated as the most acidic site [38]. In the case of l-ascorbic acid, the monoanionic form, derived from the deprotonation of the hydroxy group linked to carbon 3 [O(3)-H] (see Figure 1c), is particularly stable thanks to the delocalization of the negative charge on the oxygen atoms at positions 1 and 3. The second deprotonation occurs at higher pH values, and the corresponding constant is difficult to measure in aqueous media [45,46,47].
An analysis of the acidity constants in the literature shows major variation among the published values. Furthermore, most of the reported pKa values were acquired using alcohol and water combinations; thus, there is limited knowledge on pKa values found in aqueous solutions. Therefore, the acidic constant values determined in this work are not directly comparable with the others in literature, owing to differences in the experimental conditions (Table A1)—except in [11] and [14], which have reported values that are in excellent agreement with our data.
Only a qualitative comparison can be carried out with the values obtained in other works, which reveals acceptable agreement. The difference between those values and those determined here could be related mainly to the different temperatures rather than to the medium used.
Our results showed excellent agreement between the two experimental approaches, providing a valuable method to determine acidic constants values, in a wide range of acidity, with accuracy and in aqueous media (i.e., from pH 2 to pH 12). The acidic constants determined in this work can be considered significant values in the chemistry of these ligands, and they could be indispensable to understanding ligand behavior in biological systems. Of particular biomedical importance is the ability of these compounds to form stable complexes with biologically active metal ions. The studies in this area have focused on the enhanced curing ability in the case of vitamin C and metal ion deficiencies; on the development of therapeutic agents with potential antitumor, antibacterial, antioxidant and anti-inflammatory properties with enhanced potency with respect to the noncomplexed ligands; and on their use as synthetic models for metal containing complex biological systems [13,18,48].

Author Contributions

Conceptualization: E.F. Investigation: L.M. Resources: E.F. Writing—original draft: L.M. and E.F. Writing—review and editing: E.F. All authors have read and agreed to the published version of the manuscript.

Funding

This research received no external funding. We thank the University of Calabria.

Institutional Review Board Statement

Not applicable.

Informed Consent Statement

Not applicable.

Data Availability Statement

Data sharing not applicable.

Conflicts of Interest

The authors declare no conflict of interest.

Appendix A

Applsci 13 00776 g0a1a 550Applsci 13 00776 g0a1b 550
Figure A1. ZH as a function of −log [H+] for quercetin (a), luteolin (b) and l-ascorbic acid (c), at 37 °C and in 0.16 M NaCl. Orange circles and orange crosses refer to CL 0.5 × 10−3 M, while blue triangles refer to CL 5 × 10−3 M.
Figure A1. ZH as a function of −log [H+] for quercetin (a), luteolin (b) and l-ascorbic acid (c), at 37 °C and in 0.16 M NaCl. Orange circles and orange crosses refer to CL 0.5 × 10−3 M, while blue triangles refer to CL 5 × 10−3 M.
Applsci 13 00776 g0a1aApplsci 13 00776 g0a1b
Applsci 13 00776 g0a2 550
Figure A2. UV-vis spectra recorded for quercetin 0.025 mM (a), luteolin 0.010 mM (b) and l-ascorbic acid 0.040 mM (c), at 37 °C and in 0.16 M NaCl.
Figure A2. UV-vis spectra recorded for quercetin 0.025 mM (a), luteolin 0.010 mM (b) and l-ascorbic acid 0.040 mM (c), at 37 °C and in 0.16 M NaCl.
Applsci 13 00776 g0a2
Table
Table A1. Acidic constants of quercetin (Que), luteolin (Lut) and l-ascorbic acid (A.A.) from the literature. The uncertainties on the log values represent σ.
Table A1. Acidic constants of quercetin (Que), luteolin (Lut) and l-ascorbic acid (A.A.) from the literature. The uncertainties on the log values represent σ.
MethodMedium°Clog K*pKa1pKa2pKa3pKa4pKa5Ref.
QueElectroph.H2O25 7.1 ± 0.19.1 ± 0.111.1 ± 0.4 [42]
Colorim.H2O251.8 ± 0.16.4 ± 0.18.1 ± 0.19.0 ± 0.19.6 ± 0.111.3 ± 0.1[31]
Spectr. 6.6 ± 0.18.1 ± 0.1 11.4 ± 0.1
LutPot.0.16 M NaCl37 8.9 ± 0.1 [14]
A.A.Cond.H2O25 4.147 [45]
Pot.0.16 M NaCl37 3.86 ± 0.01 [11]
Spectr.H2O25 4.16 ± 0.0111.73 ± 0.02 [46]
Titr.H2O16–18 4.1411.43 [47]

References

  1. Beltran, J.L.; Sanli, N.; Fonrodona, G.; Barron, D.; Ozkan, G.; Barbosa, J. Spectrophotometric, potentiometric and chromatographic pKa values of polyphenolic acids in water and acetonitrile–water media. Anal. Chim. Acta 2003, 484, 253–264. [Google Scholar] [CrossRef]
  2. Erdemgil, F.Z.; Sanli, S.; Sanli, N.; Ozkan, G.; Barbosa, J.; Guiteras, J.; Beltran, J.L. Determination of pK(a) values of some hydroxylated benzoic acids in methanol-water binary mixtures by LC methodology and potentiometry. Talanta 2007, 72, 489–496. [Google Scholar] [CrossRef]
  3. Ragnar, M.; Lindgren, C.T.; Nilvebrant, N.-O. pKa-Values of guaiacyl and syringyl phenols related to lignin. J. Wood Chem. Technol. 2000, 20, 277–305. [Google Scholar] [CrossRef]
  4. Lee, S.K.; Mbwambo, Z.H.; Chung, H.S.; Luyengi, L.; Gamez, E.J.C.; Mehta, R.G.; Kinghorn, A.D.; Pezzuto, J.M. Evaluation of the antioxidant potential of natural products. Comb. Chem. High Throughput Screen. 1998, 1, 35–46. [Google Scholar] [CrossRef] [PubMed]
  5. Di Donna, L.; Bartella, L.; De Vero, L.; Gullo, M.; Giufrè, A.M.; Zappia, C.; Capocasale, M.; Poiana, M.; D’Urso, S.; Caridi, A. Vinegar production from Citrus bergamia by-products and preservation of bioactive compounds. Eur. Food Res. Technol. 2020, 246, 1981–1990. [Google Scholar] [CrossRef]
  6. Bartella, L.; Mazzotti, F.; Talarico, I.R.; De Luca, G.; Santoro, I.; Prejanò, M.; Riccioni, C.; Marino, T.; Di Donna, L. Structural Characterization of Peripolin and Study of Antioxidant Activity of HMG Flavonoids from Bergamot Fruit. Antioxidants 2022, 11, 1847. [Google Scholar] [CrossRef]
  7. Nijveldt, R.J.; van Nood, E.; van Hoorn, D.E.C.; Boelens, P.G.; van Norren, K.; van Leeuwen, P.A.M. Flavonoids: A review of probable mechanisms of action and potential applications. Am. J. Clin. Nutr. 2001, 74, 418–425. [Google Scholar] [CrossRef] [Green Version]
  8. Taverna, D.; Di Donna, L.; Mazzotti, F.; Tagarelli, A.; Napoli, A.; Furia, E.; Sindona, G. Rapid discrimination of bergamot essential oil by paper spray mass spectrometry and chemometric analysis. J. Mass Spectrom. 2016, 51, 761–767. [Google Scholar] [CrossRef]
  9. Corradini, E.; Foglia, P.; Giansanti, P.; Gubbiotti, R.; Samperi, R.; Laganà, A. Flavonoids: Chemical properties and analytical methodologies of identification and quantitation in foods and plants. Nat. Prod. Res. 2011, 25, 469–495. [Google Scholar] [CrossRef]
  10. Furia, E.; Marino, T.; Russo, N. Insights into the coordination mode of quercetin with Al(III) ion from a combined experimental and theoretical study. Dalton Trans. 2014, 43, 7269–7274. [Google Scholar] [CrossRef]
  11. Cesario, D.; Furia, E.; Mazzone, G.; Beneduci, A.; De Luca, G.; Sicilia, E. The Complexation of Al3+ and Ni2+ by l-Ascorbic Acid: An Experimental and Theoretical Investigation. J. Phys. Chem. A 2017, 121, 9773–9781. [Google Scholar] [CrossRef] [PubMed]
  12. Corrente, G.A.; Malacaria, L.; Beneduci, A.; Furia, E.; Marino, T.; Mazzone, G. Experimental and theoretical study on the coordination properties of quercetin towards aluminum(III), iron(III) and copper(II) in aqueous solution. J. Mol. Liq. 2021, 325, 115171. [Google Scholar] [CrossRef]
  13. Malacaria, L.; Corrente, G.A.; Beneduci, A.; Furia, E.; Marino, T.; Mazzone, G. A review on coordination properties of Al(III) and Fe(III) towards natural antioxidant molecules: Experimental and theoretical insights. Molecules 2021, 26, 2603. [Google Scholar] [CrossRef] [PubMed]
  14. Malacaria, L.; La Torre, C.; Furia, E.; Fazio, A.; Caroleo, M.C.; Cione, E.; Gallelli, L.; Marino, T.; Plastina, P. Aluminum(III), iron(III) and copper(II) complexes of luteolin: Stability, antioxidant, and anti-inflammatory properties. J. Mol. Liq. 2022, 345, 117895. [Google Scholar] [CrossRef]
  15. Ritacca, A.G.; Malacaria, L.; Sicilia, E.; Furia, E.; Mazzone, G. Experimental and theoretical study of the complexation of Fe3+ and Cu2+ by l-ascorbic acid in aqueous solution. J. Mol. Liq. 2022, 355, 118973. [Google Scholar] [CrossRef]
  16. Hilgers, R.; Bijlsma, J.; Malacaria, L.; Vincken, J.-P.; Furia, E.; De Bruijn, W.J.C. Transition metal cations catalyze 16O/18O exchange of catechol motifs with H218O. Org. Biomol. Chem. 2022, 20, 9093–9097. [Google Scholar] [CrossRef]
  17. Malacaria, L.; Bijlsma, J.; Hilgers, R.; De Bruijn, W.J.C.; Vincken, J.-P.; Furia, E. Insights into the complexation and oxidation of quercetin and luteolin in aqueous solutions in presence of selected metal cations. J. Mol. Liq. 2023, 369, 120840. [Google Scholar] [CrossRef]
  18. Crisponi, G.; Nurchi, V.M.; Bertolasi, V.; Remelli, M.; Faa, G. Chelating agents for human diseases related to aluminium overload. Coord. Chem. Rev. 2012, 256, 89–104. [Google Scholar] [CrossRef]
  19. Nurchi, V.M.; Crespo-Alonso, M.; Toso, L.; Lachowicz, J.I.; Crisponi, G. Chelation Therapy for Metal Intoxication: Comments from a Thermodynamic Viewpoint. Mini-Rev. Med. Chem. 2013, 13, 1541–1549. [Google Scholar] [CrossRef]
  20. Hofer, T.; Jørgensen, T.O.; Olsen, R.L. Comparison of food antioxidants and iron chelators in two cellular free radical assays: Strong protection by Luteolin. J. Agric. Food Chem. 2014, 62, 8402–8410. [Google Scholar] [CrossRef]
  21. Crisponi, G.; Dean, A.; Di Marco, V.; Lachowicz, J.I.; Nurchi, V.M.; Remelli, M.; Tapparo, A. Different approaches to the study of chelating agents for iron and aluminium overload pathologies. Anal. Bioanal. Chem. 2013, 405, 585–601. [Google Scholar] [CrossRef]
  22. Drüeke, T.B. Intestinal absorption of aluminium in renal failure. Nephrol. Dial. Transplant. 2002, 17, 13–16. [Google Scholar] [CrossRef] [Green Version]
  23. Cherrak, S.A.; Mokhtari-Soulimane, N.; Berroukeche, F.; Bensenane, B.; Cherbonnel, A.; Merzouk, H.; Elhabiri, M. In Vitro Antioxidant versus Metal Ion Chelating Properties of Flavonoids: A Structure-Activity Investigation. PLoS ONE 2016, 11, e0165575. [Google Scholar] [CrossRef] [Green Version]
  24. Selvaraj, S.; Krishnaswamy, S.; Devashya, V.; Sethuraman, S.; Krishnan, U.M. Flavonoid-Metal Ion Complexes: A Novel Class of Therapeutic Agents. Med. Res. Rev. 2014, 34, 677–702. [Google Scholar] [CrossRef] [PubMed]
  25. Kostyuk, V.A.; Potapovich, A.I.; Kostyuk, T.V.; Cherian, M.G. Metal complexes of dietary flavonoids: Evaluation of radical scavenger properties and protective activity against oxidative stress in vivo. Cell. Mol. Biol. 2007, 53, 62–69. [Google Scholar] [PubMed]
  26. Biedermann, G.; Sillén, L.G. Studies on the hydrolysis of metal ions. IV. Liquid junction potentials and constancy of activity factors in NaClO4–HClO4 ionic medium. Arkiv Kemi. 1953, 5, 425–440. [Google Scholar]
  27. Gran, G. Determination of the equivalent point in potentiometric titrations. Acta Chem. Scand. 1950, 4, 559–577. [Google Scholar] [CrossRef] [Green Version]
  28. Gran, G. Determination of the equivalence point in potentiometric titrations. Part II. Analyst 1952, 77, 661–670. [Google Scholar] [CrossRef]
  29. Beneduci, A.; Corrente, G.A.; Marino, T.; Aiello, D.; Bartella, L.; Di Donna, L.; Napoli, A.; Russo, N.; Romeo, I.; Furia, E. Insight on the chelation of aluminum(III) and iron(III) by curcumin in aqueous solution. J. Mol. Liq. 2019, 296, 111805. [Google Scholar] [CrossRef]
  30. Gans, P.; Sabatini, A.; Vacca, A. Investigation of equilibria in solution. Determination of equilibrium constants with the HYPERQUAD suite of programs. Talanta 1996, 43, 1739–1753. [Google Scholar] [CrossRef]
  31. Chebotarev, A.N.; Snigur, D.V. Study of the Acid Base Properties of Quercetin in Aqueous Solutions by Color Measurements. J. Anal. Chem. 2015, 70, 55–59. [Google Scholar] [CrossRef]
  32. Gans, P.; Sabatini, A.; Vacca, A. SUPERQUAD: An improved general program for computation of formation constants from potentiometric data. J. Chem. Soc. Dalton Trans. 1985, 1195–1200. [Google Scholar] [CrossRef]
  33. Sillén, L.G. Some Graphical Methods for Determining Equilibrium Constants. II. On “Curve-fitting” Methods for Two-variable Data. Acta Chem. Scand. 1956, 10, 186–202. [Google Scholar] [CrossRef]
  34. De Stefano, C.; Foti, C.; Giuffrè, O.; Sammartano, S. Dependence on Ionic Strength of Protonation Enthalpies of Polycarboxylate Anions in NaCl Aqueous Solution. J. Chem. Eng. Data 2001, 46, 1417–1424. [Google Scholar] [CrossRef]
  35. Zhang, C.; Korshin, G.V.; Kuznetsov, A.M.; Yan, M. Experimental and quantum-chemical study of differential absorbance spectra of environmentally relevant species: A study of quercetin deprotonation and its interactions with copper (II) ions. Sci. Total Environ. 2019, 679, 229–236. [Google Scholar] [CrossRef]
  36. Álvarez-Diduk, R.; Ramírez-Silva, M.T.; Galano, A.; Merkoçi, A. Deprotonation mechanism and acidity constants in aqueous solution of flavonols: A combined experimental and theoretical study. J. Phys. Chem. B 2013, 117, 12347–12359. [Google Scholar] [CrossRef]
  37. Musialik, M.; Kuzmicz, R.; Pawłowski, T.S.; Litwinienko, G. Acidity of Hydroxyl Groups: An Overlooked Influence on Antiradical Properties of Flavonoids. J. Org. Chem. 2009, 74, 2699–2709. [Google Scholar] [CrossRef]
  38. Amat, A.; De Angelis, F.; Sgamellotti, A.; Fantacci, S. Acid-base chemistry of Luteolin and its methyl-ether derivatives: A DFT and ab initio investigation. Chem. Phys. Lett. 2008, 462, 313–317. [Google Scholar] [CrossRef]
  39. Agrawal, P.K.; Schneider, H.-J. Deprotonation induced 13C NMR shifts in phenols and flavonoids. Tetrahedron Lett. 1983, 24, 177–180. [Google Scholar] [CrossRef]
  40. Tyukavkina, N.A.; Pogodaeva, N.N. Ultraviolet Absorption of Flavonoids VIII. Ionization Constants of Kaempferol and Quercetin. Chem. Nat. Compd. 1975, 11, 741–743. [Google Scholar] [CrossRef]
  41. Zenkevich, I.G.; Guschina, S.V. Determination of Dissociation Constants of Species Oxidizable in Aqueous Solution by Air Oxygen on an Example of Quercetin. J. Anal. Chem. 2010, 65, 371–375. [Google Scholar] [CrossRef]
  42. Herrero-Martínez, J.M.; Sanmartin, M.; Rosés, M.; Bosch, E.; Ràfols, C. Determination of dissociation constants of flavonoids by capillary electrophoresis. Electrophoresis 2005, 26, 1886–1895. [Google Scholar] [CrossRef] [PubMed]
  43. Jovanovic, S.V.; Steenken, S.; Tosic, M.; Marjanovic, B.; Simic, M.G. Flavonoids as Antioxidants. J. Am. Chem. Soc. 1994, 116, 4846–4851. [Google Scholar] [CrossRef]
  44. Herrero-Martínez, J.M.; Repollés, C.; Bosch, E.; Rosés, M.; Ràfols, C. Potentiometric Determination of Aqueous Dissociation Constants of Flavonols Sparingly Soluble in Water. Talanta 2008, 74, 1008–1013. [Google Scholar] [CrossRef] [PubMed]
  45. Rimpapa, Z.; Pleho-Kapić, A.; Galijašević, S.; Šapčanin, A.; Korać, F. Change in Acidity of l-Ascorbic Acid in the Mixed Solvent DMSO—Water Followed by Conductometric Determination of Dissociation Constants. Bull. Chem. Technol. Bosnia. Herzeg. 2013, 40, 35–38. [Google Scholar]
  46. Seok, Y.-J.; Kap-Seok, Y.; Sa-Ouk, K. A simple spectrophotometric determination of dissociation constants of organic compounds. Anal. Chim. Acta 1995, 306, 351–356. [Google Scholar] [CrossRef]
  47. Birch, T.W.; Harris, L.J. The titration curve and dissociation constants of vitamin C. Biochem. J. 1933, 27, 595. [Google Scholar]
  48. Zumreoglu-Karan, B. The coordination chemistry of Vitamin C: An overview. Coord. Chem. Rev. 2006, 250, 2295–2307. [Google Scholar] [CrossRef]
Applsci 13 00776 g001 550
Figure 1. Chemical structure of the investigated natural antioxidants: quercetin (a), luteolin (b) and l-ascorbic acid (c).
Figure 1. Chemical structure of the investigated natural antioxidants: quercetin (a), luteolin (b) and l-ascorbic acid (c).
Applsci 13 00776 g001
Table
Table 1. Survey of the log K* and log Km values, according to general equilibria 3 and 4, by numerical methods. The uncertainties represent 3 σ.
Table 1. Survey of the log K* and log Km values, according to general equilibria 3 and 4, by numerical methods. The uncertainties represent 3 σ.
log K*log Ka1log Ka2log Ka3log Ka4log Ka5
QuercetinSuperquad
2.00 ± 0.06−8.29 ± 0.03−8.61 ± 0.01−9.5 ± 0.1−9.7 ± 0.3−10.4 ± 0.3
Hyperquad
/−8.5 ± 0.6−9.0 ± 0.3−9.5 ± 0.6−10.0 ± 0.3−10.5 ± 0.9
LuteolinSuperquad
2.3 ± 0.1−8.29 ± 0.03−8.6 ± 0.1−8.8 ± 0.3−9.3 ± 0.3
Hyperquad
/−8.1 ± 0.3−8.80 ± 0.05−9.6 ± 0.6−9.8 ± 0.6
l-Ascorbic Acidlog K* log Ka1 log Ka2
Superquad
1.2 ± 0.1-−3.86 ± 0.03 /
Hyperquad
1.16 ± 0.05 −3.75 ± 0.06 −10.6 ± 0.1
Disclaimer/Publisher’s Note: The statements, opinions and data contained in all publications are solely those of the individual author(s) and contributor(s) and not of MDPI and/or the editor(s). MDPI and/or the editor(s) disclaim responsibility for any injury to people or property resulting from any ideas, methods, instructions or products referred to in the content.

Share and Cite

MDPI and ACS Style

Malacaria, L.; Furia, E. Wide pH Range Potentiometric and Spectrophotometric Investigation into the Acidic Constants of Quercetin, Luteolin and l-Ascorbic Acid in Aqueous Media. Appl. Sci. 2023, 13, 776. https://doi.org/10.3390/app13020776

AMA Style

Malacaria L, Furia E. Wide pH Range Potentiometric and Spectrophotometric Investigation into the Acidic Constants of Quercetin, Luteolin and l-Ascorbic Acid in Aqueous Media. Applied Sciences. 2023; 13(2):776. https://doi.org/10.3390/app13020776

Chicago/Turabian Style

Malacaria, Luana, and Emilia Furia. 2023. "Wide pH Range Potentiometric and Spectrophotometric Investigation into the Acidic Constants of Quercetin, Luteolin and l-Ascorbic Acid in Aqueous Media" Applied Sciences 13, no. 2: 776. https://doi.org/10.3390/app13020776

APA Style

Malacaria, L., & Furia, E. (2023). Wide pH Range Potentiometric and Spectrophotometric Investigation into the Acidic Constants of Quercetin, Luteolin and l-Ascorbic Acid in Aqueous Media. Applied Sciences, 13(2), 776. https://doi.org/10.3390/app13020776

Note that from the first issue of 2016, this journal uses article numbers instead of page numbers. See further details here.

Article Metrics

Back to TopTop