INTRODUCTION

Iron is a vital trace element that regulates the processes of respiration, cell growth, metabolism, DNA synthesis and repair [1–3]. In particular, iron(III) in the form of heme is a component of cytochrome P450, whose isoenzymes are involved in the first phase of the metabolism of a large number of exogenous and endogenous compounds and in the biosynthesis of steroid hormones, cholesterol, bile acids, and prostanoids [4]. Obviously, iron(III) acts as an active binding site for various bioligands containing donor atoms of nitrogen, oxygen, and sulfur in their structure. Typical representatives of the latter are barbituric acids, belonging to the class of exogenous origin ligands.

Barbituric acid (C4H4N2O3, H2BA) and 2-thiobarbituric acid (C4H4N2O2S, H2TBA) are the parent molecules of a large group of drugs with anesthetic, antisclerotic, anticonvulsant, and sedative effects [5, 6]. Some members of the class demonstrate antibacterial, antifungal, anti-inflammatory, antiviral, and anticancer activity [5–10]. In analytical practice, H2BA and H2TBA are widely used in the colorimetric determination of cyanides (König reaction) [11] and malondialdehyde (MDA) [12], respectively. In aqueous solutions, acids (R = O for H2BA and R = S for H2TBA) exist predominantly in the keto (left) and enol (right) forms [7]:

The presence of electron-donating nitrogen, oxygen, and sulfur atoms in the imide >N–H, carbonyl >C=O, and thiocarbonyl >C=S groups determines the ability of H2BA and H2TBA to act as mono- or bidentate ligands. The acidic properties are due to the hydrogen atoms of the hydroxyl –OH (given the possible enolization of the >C=O group) and imide groups (both forms, positions 1 and 6).

There are data in the literature on the composition and properties of iron(III) compounds with H2BA [13, 14] and H2TBA [7, 15]. Due to the formation of intensely colored complexes, soluble iron(III) salts are used both in the quantitative assessment of the content of these barbituric acids (pH 2–4) [13, 16] and in determining the authenticity of their derivatives in pharmaceutical chemistry [17]. Morelli [18] suggested using H2TBA as an analytical reagent for Cu2+ and Fe3+ ions in their joint determination without preliminary separation. Using the method of molar ratios at various pH values, the 1 : 2 stoichiometric composition of the dominant complexes in the FeCl3–H2TBA system (λmax = 375 nm) was established. Nevertheless, complexation between the studied components in an aqueous solution are practically not studied to date, especially this concerns the quantitative characteristics of equilibria (Table 1).

Table 1. Results of the stability constants determination for complex particles of iron(II, III) with barbituric and 2-thiobarbituric anions in an aqueous solution by spectrophotometry and pH-metry
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In 1964, Wills [20] noticed that the spectrophotometric determination of MDA by the thiobarbiturate method in the presence of iron(III) leads to erroneous results because of a significant increase in absorption (λ = 535 nm). The author suggested that the reason was the formation of iron(III) 2-thiobarbiturate complexes. Later, Zhou and colleague [21] showed that the previously described interfering effect of Fe3+ ions is indeed due to the formation of complex particles between iron(III) and H2TBA, which have pronounced absorption at 532 nm or fluorescence at 553 nm (characteristic maxima for the MDA–H2TBA adduct). The authors give the probable stoichiometric composition of the complex as 1 : 3 (precipitation was observed). Thus, the presence of iron, given its biogenic origin, can lead to false positive results of biochemical tests when quantifying MDA, which is a marker of oxidative stress, by the thiobarbiturate method. According to the described spectrophotometric procedures [12, 22–24], the determination of MDA in the form of an adduct with H2TBA is carried out in a strongly acid medium using orthophosphoric and glacial acetic acids, as well as its chlorine-substituted derivatives.

The aim of this work is to study complexation in the iron(III)–barbituric (2-thiobarbituric) acid system, including determination of the composition and thermodynamic stability of complex particles in an aqueous solution in an acidic medium.

EXPERIMENTAL

Reagents and equipment. Barbituric (Reakhim, Russia) and 2-thiobarbituric (Diam, Russia) acids (pure for analysis grade, Russian State Standard), iron(III) chloride hexahydrate (Reakhim, Russia, chemically pure grade), sodium hydroxide (Khimreaktiv, Russia, pure for analysis grade), sodium chloride (Base No. 1 Khimreaktivov, Russia, chemically pure grade), potassium dihydroorthophosphate (Base No. 1 Khimreaktivov, Russia, chemically pure grade), potassium fluoride dihydrate (Reakhim, Russia, pure for analysis grade), orthophosphoric (Khimprom, Russia, chemically pure grade) and hydrochloric (Sigma Tech, Russia, special quality grade) acids, sulfosalicylic acid dihydrate (Base No. 1 Khimreaktivov, Russia, chemically pure grade), Trilon B (standard titer, Reakhim, Russia) were used as starting reagents. The working solutions were maintained at a constant ionic strength I = 0.1 (NaCl) and temperature t = 20°C. The concentration of H2L ligands was preliminarily refined by pH-metric titration with an alkali solution; iron(III) was refined by direct complexometric titration with a Trilon B solution using sulfosalicylic acid as an indicator according to a well-known method [25]. To prevent the hydrolysis process, the iron(III) chloride solution was acidified with hydrochloric acid. In the experiments, we utilized ultrapure water (type I) obtained using a specialized Millipore Direct-Q 5 UV water purification system (Merck, USA).

Spectrophotometric measurements in the UV and visible spectrum regions were carried out on a Cary-50 spectrophotometer (Varian, USA) in quartz cuvettes (l = 10 mm, against water). pH was measured using a 673 pH meter (Analitpribor, Georgia) on a concentration scale. For that, glass electrode was preliminarily calibrated against (H, Na)Cl solutions with different HCl concentrations (I = 0.1). pH-metric titration in determining the protonation constants of ligand anions and the stability constants of complex particles was carried out with a solution of preliminarily standardized carbonate-free alkali (I = 0.1) with continuous stirring and passage of purified nitrogen through the titrated solution.

Choice of ligand protonation constants. The protonation constants of singly (HL–) and doubly charged (L2–) anions of the studied ligands were used in the calculations, characterizing the following equilibria:

$$\begin{gathered} {\text{H}}{{{\text{L}}}^{-}} + {\text{ }}{{{\text{H}}}^{ + }}~\xrightarrow{{{{\varkappa }_{2}}}}{{{\text{H}}}_{{\text{2}}}}{\text{L}}, \\ {\text{log}}\,{{\varkappa }_{2}} = {\text{ }}3.69{\text{ }}\left( {{{{\text{H}}}_{{\text{2}}}}{\text{BA}}} \right),{\text{ }}2.30{\text{ }}\left( {{{{\text{H}}}_{{\text{2}}}}{\text{TBA}}} \right),~ \\ \end{gathered} $$
(1)
$$\begin{gathered} {{{\text{L}}}^{2}}^{-} + {\text{ }}{{{\text{H}}}^{ + }}~\xrightarrow{{{{B}_{1}}}}{\text{H}}{{{\text{L}}}^{-}},{\text{ }} \\ {\text{log}}\,{{B}_{1}} = 11.90\left( {{{{\text{H}}}_{{\text{2}}}}{\text{BA}}} \right),\,\,10.55\left( {{{{\text{H}}}_{{\text{2}}}}{\text{TBA}}} \right), \\ \end{gathered} $$
(2)
$$\begin{gathered} {{{\text{L}}}^{2}}^{-} + 2{{{\text{H}}}^{ + }}~\xrightarrow{{{{B}_{2}}}}{{{\text{H}}}_{{\text{2}}}}{\text{L,}} \\ {\text{log}}\,{{B}_{2}} = {\text{ }}15.59\left( {{{{\text{H}}}_{{\text{2}}}}{\text{BA}}} \right),\,\,12.85\left( {{{{\text{H}}}_{{\text{2}}}}{\text{TBA}}} \right),~~~~~~~~~~~ \\ \end{gathered} $$
(3)

where ϰ2 is the stepwise protonation constant; B1 and B2 are the overall protonation constants of the L2– ion.

The stepwise protonation constants ϰ2 (in log units) of the ligands were preliminarily determined by pH-metry (I = 0.1, T = 20°C): 3.69 ± 0.02 (H2BA), 2.30 ± 0.01 (H2TBA). The overall protonation constants for the first step B1 (in log units) were taken from [19, 26]: 11.90 ± 0.04 for H2BA (I = 0.1 (KNO3), T = 25°C) and 10.55 ± 0.10 for H2TBA (I = 0.1 (NaClO4), T = 18°C). The overall protonation constants for the second step B2 are obtained from the equation: log B2 = log B1 + logϰ2. The temperature difference was not taken into account, since it was insignificant (according to [19], an increase in temperature by 13°C led to a change in the value of logB1 by only 0.11 units).

Determination of the composition and stability of complex particles. The composition and stability of the complex particles formed in the studied systems under specified conditions were determined from the data of spectrophotometric (isomolar series, Benesi–Hildebrand, D–pH methods [27, 28]) and pH metric measurements in the pH range of 1.3–3.3. The effective wavelength, where the observed absorption is maximum for the complex and minimum for the free metal and ligand (DL = 0) was taken to be 440 or 540 nm for the FeCl3–H2BA system and 490 or 540 nm for the FeCl3–H2TBA system. Before absorbance measurements, the mixtures were kept for 1 min until a nearly constant Di value was established.

Processing of measurement and calculation results. Calculation of protonation constants (“Acid 1”) and stability constants of complex particles according to spectrophotometry (“D–pH”) and pH-metry (“Bjerrum 1”), construction of species distribution diagrams of the studied ligands (“Acid Yield”), as well as statistical processing of the obtained results (“Delta”) was performed using the programs indicated in brackets [28]. When calculating the overall stability constants of complex particles, we took into account the possible protonation of the ligand (Eqs. (1)–(3)) and hydrolysis of the metal ion in the first step (log Kh1 = –2.84 [28]). In the case of the D–pH series, the complexation of Fe3+ ions with supporting electrolyte anions (Cl–) was also given. For each studied system, at least two parallel experiments were performed.

Iron(III) masking in the presence of 2-thiobarbiturate anion. When assessing the possibility of iron(III) masking in the presence of 2-thiobarbiturate anion, a number of well-known masking agents were used: orthophosphate, fluoride, and ethylenediaminetetraacetate anions. The reference samples were aqueous solutions of 1.68 × 10–3 M FeCl3 (solution 1, pH 1.35) and FeCl3–H2TBA (1 : 1, solution 2, СFe = CL = 1.68 × 10–3 M, pH 1.36). Next, a stoichiometric amount of an aqueous solution of Trilon B (solution 3, pH 1.53) was added to the resulting mixture; solid KF (solution 4, pH 1.89) and KH2PO4 (solution 5, pH 1.77) were also added at the tip of a spatula. Then the solutions were mixed, left for 10 min, and their color was fixed (Fig. 1). In all cases, the constant ionic strength I = 0.1 was maintained by adding a mixture of (H, Na)Cl at T = 20°C. Instead of KH2PO4, 10 μL of concentrated H3PO4 (Vtotal = 6 mL) was also added to solution 5. All test solutions were stable for at least 24 h.

Fig. 1.
figure 1

Visual representation of solutions used in assessing the possibility of iron(III) masking in the presence of H2TBA (CL = CFe = Сmixture = 1.68 × 10–3 М, Vtotal = 6 mL, I ~ 0.1, t = 20°C): (1) FeCl3 (pH 1.35), (2) FeCl3–H2TBA (pH 1.36), (3) FeCl3–H2TBA + Na2H2Edta (Cadditive = 1.68 × 10–3 М, pH 1.53), (4) FeCl3–H2TBA + KF(solid) (pH 1.89), (5) FeCl3–H2TBA + KH2PO4(solid) (pH 1.77).

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RESULTS AND DISCUSSION

Composition and Stability of Complex Particles

When aqueous solutions of FeCl3 and H2L were mixed, the resulting solution turned red-brown accompanied by a significant increase in absorption in the UV and visible (DL = 0) spectrum regions (Figs. 2a and 2b). The electronic absorption spectra of H2L solutions in the UV range (Figs. 2a and 2b) contain pronounced maxima (Table 2) due to intraligand transitions π → π* and n → π* in the heterocyclic fragments of the considered acids and their functional groups [14], as well as those undergoing a hypsochromic shift in the absorption spectra of the complexes. The absorption band of the FeCl3–H2BA and FeCl3–H2TBA systems in the visible range (Figs. 2a and 2b) forms a pronounced shoulder emanating from the UV region and may belong to a charge transfer complex. According to Ref. [13], the spectra of the FeCl3–H2BA system also contained a charge transfer band in the region of 430–480 nm. The above facts unambiguously indicate the occurrence of complexation processes in the studied systems. This is also indirectly confirmed by the fact that the titration curves of the FeCl3–H2L systems are located significantly lower than the titration curves of the acids themselves, since the Fe3+ ion promotes the displacement of the ligand proton [13].

Fig. 2.
figure 2

Electronic absorption spectra (I ~ 0.1, pH ~ 3.0, T = 20°C) of FeCl3–H2L (1), H2L (2) and FeCl3 (3) aqueous solutions in UV (CL = CFe = Сmixture = 3.0 × 10–5 М) and visible region (CL = CFe = Сmixture = 2.5 × 10–2 М), where H2L = (a) H2BA or (b) H2TBA.

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Table 2. Results of UV spectrophotometric study of FeCl3–H2L systems and free ligands in aqueous solution (CFe = CL = 3.0 × 10–5 М, I = 0.1, pH ~ 3.0, T = 20°C, l = 10 mm)
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The determination of the composition and thermodynamic stability of iron(III) complex particles with H2BA and H2TBA anions was preceded by a study of the acid–base properties of the ligands by pH-metry. The obtained values of the logarithms of the stepwise protonation constants of the anions BA2– (3.69 ± 0.02) and TBA2– (2.30 ± 0.01) at I = 0.1 (NaCl) and T = 20°C satisfactorily agree with the currently available literature data [7, 13, 19 , 26].

The study of the FeCl3–H2L systems using the isomolar series (Fig. 3) and D–pH (M : L ~ 1 : 1) showed dominance of the 1 : 1 complex particles containing the protonated form of the ligand under experimental conditions (pH 1.3–2.4, I = 0.1, T = 20°C). At the same time, according to the results of pH-metry (pH 2.6–3.1), the simultaneous presence of protonated [FeHL]2+ and deprotonated [FeL]+ complexes was established. Thus, given the species distribution diagrams of H2L, the formation of barbiturate and 2-thiobarbiturate iron(III) complexes in an aqueous solution can be represented by the following equilibria:

$${\text{F}}{{{\text{e}}}^{{3 + }}} + {{{\text{H}}}_{{\text{2}}}}{\text{L}}\overset {{{K}_{1}}} \longleftrightarrow \,\,{{[{\text{FeHL}}]}^{{2 + }}} + {{{\text{H}}}^{ + }},$$
(4)
$${\text{F}}{{{\text{e}}}^{{3 + }}} + {{{\text{H}}}_{{\text{2}}}}{\text{L}}\overset {{{K}_{2}}} \longleftrightarrow \,\,{{[{\text{FeL}}]}^{ + }} + 2{{{\text{H}}}^{ + }},$$
(5)
$${\text{F}}{{{\text{e}}}^{{3 + }}} + {\text{H}}{{{\text{L}}}^{ - }}\overset {{{K}_{3}}} \longleftrightarrow \,\,{{[{\text{FeL}}]}^{ + }} + {{{\text{H}}}^{ + }},$$
(6)
$${\text{F}}{{{\text{e}}}^{{3 + }}} + {\text{H}}{{{\text{L}}}^{ - }}\overset {{{\beta }_{{11}}}} \longleftrightarrow \,\,{{[{\text{FeHL}}]}^{{2 + }}},$$
(7)
$${\text{F}}{{{\text{e}}}^{{3 + }}} + {{{\text{L}}}^{{2 - }}}\overset {{{\beta }_{1}}} \longleftrightarrow \,\,{{[{\text{FeL}}]}^{ + }},$$
(8)
$${{{\text{[FeL]}}}^{ + }} + {{{\text{H}}}^{ + }}\overset k \longleftrightarrow \,\,{{[{\text{FeHL}}]}^{{2 + }}},$$
(9)
$${\text{F}}{{{\text{e}}}^{{3 + }}} + {{{\text{H}}}^{ + }} + {{{\text{L}}}^{{2 - }}}\overset R \longleftrightarrow \,\,{{[{\text{FeHL}}]}^{{2 + }}}.$$
(10)
Fig. 3.
figure 3

Dependence of the absorption D on the molar content of the ligand NL in a series of isomolar solutions of the system FeCl3–H2TBA: Di (1), ΔDi (2), DFe, i (3) (i is the experimental point number, λef = 540 nm, l = 10 mm, \(C_{{\text{L}}}^{^\circ }\) = \(C_{{\text{M}}}^{^\circ }\) = 6.00 × 10–3 М, ΔDmax = 0.144, Vtotal = 6 mL, pH ~ 2.5, I = 0.1, T = 20°C).

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Based on the results of isomolar series, using regions with a molar ratio of components close to 1 : 1, we calculated the stability constants of complex particles [FeHBA]2+ (log β11 = 3.29 ± 0.08, pH ~ 1.9) and [FeTBA]+ (log β1 = 10.83 ± 0.12, pH ~ 2.5). The stability constants of acidic [FeHBA]2+ (log β11 = 3.71 ± 0.19, pH ~ 2.1) and deprotonated [FeBA]+ (log β1 = 12.13 ± 0.23, pH ~ 3.3) complexes were also determined by Benesi-Hildebrand method at CFe \( \gg \) CL and CL \( \gg \) CFe. In addition, processing of the experimental data obtained by the D–pH method for the FeCl3–H2BA system (pH 1.3–2.4) led to the following results: n ≈ 1 and log R = 15.55 ± 0.20, where n is the number of protons in the complex, R is equilibrium constant (10). It follows from Eq. (2), (7), and (10) that log β11 = log R – log B1 = 15.55 – 11.90 = 3.65 (s2 = 3.02 × 10–3). For the FeCl3–H2TBA system (pH 1.6–2.4), we have n ≈ 1 and log R = 13.00 ± 0.90; thus, log β11 = 13.00 – 10.55 = 2.45 (s2 = 4.23 × 10–3).

These methods assume the predominance of only a single complex in the equilibrium mixture. However, the expediency of these assumptions (taking into account only complex particles of the same type) is confirmed by pH-metry data, which makes it possible to take into account the simultaneous presence of several complex particles of different composition. Thus, the processing of pH-metric data for the FeCl3–H2BA system (pH 2.6–3.0; Table 3) led to the following results: log β1 = 12.30 ± 0.03, log β11 = log Ropt – log B1 = 15.20 – 11.90 = 3.30 (s2 = 1.03 × 10–2). For the system FeCl3–H2TBA (pH 2.6–3.1) we have the following: log β1 = 11.27 ± 0.03, log β11 = log Ropt – log B1 = 13.47 – 10.55 = 2.92 (s2 = 4.23 × 10–3). The values of the stability constants of [FeHL]2+ and [FeL]+ complex particles calculated by different methods are in satisfactory agreement both with each other and with the currently known literature data [13] (Table 1), which indicates the correctness of the chosen calculation model and the processes embedded in it.

Table 3. pH-metry data and calculation results of the [FeBA]+ stability constant with the simultaneous presence of [FeHBA]2+ (\(V_{{\text{M}}}^{^\circ }\) = \(V_{{\text{L}}}^{^\circ }\) = 5.00 mL, \(C_{{\text{L}}}^{^\circ }\) = 1.04 × 10–2 М, \(C_{{\text{M}}}^{^\circ }\) = 1.00 × 10–2 М, СНCl = 1.142 × 10–3 М, CNaOH = 1.68 × 10–2 М, log Kh1 = –2.84, log B1 = 11.90, log B2 = 15.59, I = 0.1, T = 20°C)
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When comparing the values of β11 and β1 (Table 1), it can be seen that [FeHL]2+ are less stable than [FeL]+. This is due to the lower denticity of the HL– ion compared to L2– one, since in the first case the nitrogen atom of the imide group is blocked by a proton and is not able to participate in the coordination by the metal ion. According to the Pearson HSAB concept and the Ahrland–Chatt–Davies classification [29], the Fe3+ cation, which has a half-filled d-sublevel and has a charge distribution close to spherical, is a typical hard acid (class “A” cation) and forms the most stable complexes with hard bases, which are ligands containing donor oxygen and nitrogen atoms. It is logical that the HL– anion acts as a monodentate ligand, being coordinated by the Fe3+ ion via the deprotonated oxygen atom of the hydroxo group (see earlier keto-enol equilibrium), the proton of which is displaced first. The L2– anion is probably coordinated via donor oxygen and nitrogen atoms.

The obtained values of β11 for [FeHTBA]2+ and [FeHBA]2+, β1 for [FeTBA]+ and [FeBA]+ correlate with each other, as well as with the value of β1 of iron(II) 2-thiobarbiturate complex [FeTBA]0 [19] (Table 1). From simple electrostatic concepts for cations with a spherical charge distribution (or close to it), it follows that the complex is more stable when it includes ions with a larger charge and a smaller size. A convenient criterion for assessing the complexing ability of ions is the value of their ionic potential φ (φ = zi/ri) [30, 31]: the complex with the higher φ value (with the same composition of complexes) is more stable. Indeed, the higher stability of the iron(III) 2-thiobarbiturate complex [FeTBA]+ compared to that of iron(II) [FeTBA]0 (Table 1) is associated with a change in the electrostatic characteristics of the cations when passing from Fe2+ (d6, rи = 0.92 Å, zи = +2, φ = 2.17) to Fe3+ (d5, rи = 0.79 Å, zи = +3, φ = 3.80). As a consequence, a stronger electrostatic M–L interaction occurs in the [FeTBA]+ complex than in the [FeTBA]0 particle. In addition, a ligand with a higher basicity forms more stable complexes: the Fe3+ ion, like the proton, preferentially binds to the more basic ligand. Thus, when comparing the stability of 2-thiobarbiturate (log B1 = 10.55, log B2 = 12.85, log ϰ2 = 2.30) and barbiturate (log B1 = 11.90, log B2 = 15.59, log ϰ2 = 3.69) complexes (Table 1), it can be seen that the anions HBA– and BA2– are more basic than HTBA– and TBA2– ones. As a result, barbiturate complexes are an order of magnitude more stable than 2-thiobarbiturate analogs.

Iron(III) Masking in the Presence of 2-Thiobarbiturate Anion

It is known that biogenic metal ions present in the biological fluids of animals and humans can have a significant interfering effect on the results of various biochemical studies. In particular, the interfering effect of Fe3+ ions in studying the lipid peroxidation by the thiobarbiturate method is due to their interaction with H2TBA and the formation of an intensely colored complex [20, 21].

It was established in Ref. [20] that the addition of various chelating agents (trilon B, gelatin, ovalbumin, diluted horse serum) to a mixture of MDA and H2TBA in the presence of iron(III) leads to a noticeable decrease in the absorption of the solution. The observed phenomenon can be explained by the binding of iron(III) into a more stable and less colored complex compared to the 2-thiobarbiturate one. Thus, for [FeEdta]– (Edta4– is the anion of ethylenediaminetetraacetic acid H4Edta), log β1 = 25.10 (I = 0.1 (NaClO4), T = 20°C) [32], which is much higher than the corresponding value for [FeTBA]+ (Table 1). Proteins contain a large number of N- and O-donor functional groups and are able to bind the Fe3+ ion into strong chelate complexes. On the other hand, the addition of sodium citrate (Na3Cit), glycine (HGly), alanine (HAla), phenylalanine (HPhe) or glutamic acid (H2Glu) does not affect the absorption of the mixture. Apparently, this is due to the fact that the forming complex particles have less or close stability compared to [FeTBA]+. For instance, for [FeGly]2+ log β1 = 8.57 (I = 0.5 (NaNO3), T = 25°C) [33]; for [FeAla]2+ log β1 = 8.96 (I = 0.5 (NaNO3), T = 25°C) [34]; for [FeGlu]+ log β1 = 11.81 (I = 0.5 (NaNO3), T = 25°C) [35]; for [FeCit]0 log β1 = 11.21 (I = 0.1 (KNO3), T = 25°C) [36].

It follows from the above that it is logical to use Trilon B to eliminate the interfering effect of iron(III) in the determination of MDA by the thiobarbiturate method in a strongly acidic medium. The forming [FeEdta]– complex is more stable than [FeTBA]+ one and does not significantly absorb in the region of 530–540 nm (the interval of the characteristic maximum for the MDA–H2TBA adduct) [37], which makes it possible to avoid possible spectral noise. Fluoride and orthophosphate ions are often used in analytical practice to mask Fe3+ ones. Since the stability of iron(III) fluoride complexes (log β3 = 12.53; I = 0.1 (KNO3), T = 25°C [38]) is higher than that of [FeTBA]+, fluoride ions can also be recommended for masking Fe3+ ones. Iron(III) orthophosphate protonated complexes, e.g., [FeHPO4]+ (log β11 = 8.59; I = 0.15 (NaCl), T = 25°C [39]), are less resistant to [FeTBA]+. In this regard, orthophosphate ions can be used as a masking agent in the region of dominance of the [FeHTBA]2+ complex at pH < 2 (Fig. 4).

Fig. 4.
figure 4

Species distribution diagram of iron(III) 2-thiobarbiturate complexes in an aqueous solution (CL = 0.01 М, log B1 = 10.55, log B2 = 12.85, log Kh1 = –2.84, log β11 = 2.69, log β1 = 11.05): [Fe3+] (α0), [FeHTBA2+] (α11), [FeTBA+] (α1).

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The assumptions made above based on a comparison of the stability constants of iron(III) complexes agree with the experimental data. As can be seen from Fig. 1, solution 4 with KF and solution 5 with KH2PO4 were colorless compared to the colored reference solution 2. Adding 10 µl of concentrated H3PO4 instead of KH2PO4 also causes the solution to become colorless. Thus, H3PO4 can be used to create an acidic medium (pH ~ 2–3 [23]), which is necessary for the reaction between MDA and H2TBA, and to mask Fe3+ ions. When Trilon B solution (solution 3) is added, a yellow color appears (wavelength range 565–590 nm), which, as expected, should not affect the absorption of the MDA–H2TBA adduct [37].

CONCLUSIONS

Using spectrophotometry and pH-metry in FeCl3–H2L–H2O systems (pH 1.3–3.3, I = 0.1, T = 20°C), where H2L is barbituric or 2-thiobarbituric acid, the formation of complex particles was shown, the metal : ligand ratio was established, the forms of the ligand in the formed complexes ([FeHL]+ and [FeL]+) and their stability constants were determined. It was shown that when studying the state of lipid peroxidation, iron(III) as an interfering endogenous component can be masked by the addition of Trilon B, KF, and KH2PO4 (or H3PO4).