SUBSTANTIATION FOR OXYWATER ZWITTERIONS AND SINGLET OXYGEN ATOMS GENERATION FROM HYDROGEN PEROXIDE MOLECULES IN AQUEOUS SOLUTIONS

Hydrogen peroxide is widely used as an oxidant. The results of thermodynamic calculations in-dicate the impossibility of spontaneous generation of hydroxyl and hydroperoxyl radicals from hydrogen peroxide in aqueous solutions. Hydrogen peroxide spontaneously decomposes in ferrous, ferric, and cupric Fenton reaction systems. Ferric xylenol orange and ferric pyridoxine complexes are oxidized rapidly and spontaneously by this oxidant. Hydrogen peroxide in aqueous solutions spontaneously oxidizes the sulfur atoms of hyposulfite anions and benzylpenicillin molecules. Thus, a hydrogen peroxide molecule generates another intermediate that differs from hydroxyl and hydroperoxyl radicals. Theoretical modeling shows that hydrogen peroxide can participate in the proton transfer reactions. Its isomerization to oxywater zwitterion with subsequent oxywater intramolecular disproportionation is a process that is very suitable for explaining all events of hydrogen peroxide decomposition and oxidative reactivity in aqueous systems. The oxywater zwitterion is a bipolar ion in which the opposite charges are localized on neighboring oxygen atoms. This determines the displacement of electron density from the negatively charged atom to the positively charged atom. As a result, the interoxygen bond heterolytically dissociates with liberation of a water molecule and formation of an oxygen atom (oxene) in concept for explanation of hydroperoxide monooxygen and dioxygen oxidative functionalization mechanisms in organic synthesis.


Introduction
Hydrogen peroxide HOOH is widely used as an oxidant for medical disinfection, in advanced oxidation processes [1], in delignification processes [2], in organic synthesis for monooxygen and dioxygen functionalization of substrates [3]. Hydrogen peroxide is a redox-signaling molecule in biological systems [4], and it is a facilitator of pathological oxidative stress via endogenous Fenton reactions [5]. There is a widespread opinion that hydrogen peroxide directly generates free radicals (hydroxyl HO • and hydroperoxyl HOO • ) through one-electron redox reactions [1,3,5] (2) Equation (1) is used for explanation of divalent (ferrous) iron and monovalent (cuprous) copper interaction with hydrogen peroxide [1,3,5] (4) Equation (2) is used for explanation of trivalent (ferric) iron and divalent (cupric) copper interaction with hydrogen peroxide [1,3,5]:  [1,3,5], but these schemes are not proved [3,[6][7][8]. The summation 1+2 or 3+5 or 4+6 gives Equation (7) of one-electron hydrogen peroxide dismutation with simultaneous generation of hydroxyl and hydroperoxyl radicals: Earlier, Chumakov A.A. et al. [9] used Equation (7) as the probable first step of non-catalytic (thermal or photochemical) hydrogen peroxide decomposition. It was shown that one hydrogen atom transfer led to increase in Gibbs free energy +39.9 kJ/mol in the gas phase. There was an assumption for endergonic activation of two hydrogen peroxide molecules associate (Fig. 1  The purpose of our present study is argumentation against the widespread view that hydrogen peroxide is a direct precursor of hydroxyl and hydroperoxyl radicals, as well as presentation of an oxywater-oxene conception for hydrogen peroxide decomposition kinetics and oxidative reactivity in aqueous systems.

I. Results
We carried out a complex study using thermodynamic calculations, performing reactions such as hydrogen peroxide decomposition in Fenton systems, model substrates oxidation (hyposulfite, benzylpenicillin, ferric xylenol orange chelates, ferric pyridoxine chelates), and electro-Fenton reactions. Besides, we used theoretical modeling.
I-1. Thermodynamic analysis. We carried out thermodynamic analysis of radical-generating reactions (Equations (3)-(7) for aqueous phase, using reference values [10] of standard thermodynamic functions of reagents (rg) and products (prd) of reactions (Table 1)   The calculated standard thermodynamic functions of reactions 3-7 are presented in Table 2.

I-2. Hydrogen peroxide decomposition in Fenton systems.
We used a digital gas volume and temperature USB-detector for recording molecular dioxygen, produced during hydrogen peroxide decomposition in ferrous, ferric, and cupric Fenton reaction systems (Fig. 2 (Fig. 2).
The ferric and cupric systems had no changes in catalyst condition after the completion of hydrogen peroxide decomposition, that is, the ferric system looked like the initial yellow ferric chloride aqueous solution and the cupric system looked like the initial blue-green cupric chloride aqueous solution. In contrast, ferrous catalyst obviously changed its condition. The initial ferrous sulfate aqueous solution was transparent and colorless. Its interaction with hydrogen peroxide resulted in precipitation of iron(III) oxide-hydroxide with rusty color. I-4. Oxidation of benzylpenicillin. The benzylpenicillin sodium salt was dissolved in aqueous hydrogen peroxide solution without any Fenton catalyst addition. The system was protected from thermal and photochemical activation. As a result, we observed the colloid solution formation, and hydrogen peroxide disproportionation with the gas-phase molecular oxygen liberation.
The NMR-spectroscopy data (Section I-7) was indicative of S-oxidation of the sulfide fragment. I-5. Oxidation of ferric xylenol orange chelates. The xylenol orange aqueous solution 5•10 -5 mol/L has yellow color and absorbs visible light with λ max =430 nm. When the equimolar amount of iron(III) chloride was added, violet color with λ max =575 nm appeared. The violet color was stable for many months. When hydrogen peroxide was added, the discoloration of solution occured very quickly within a few minutes. The final solution had no absorption maximums in visible region and it looked like pale yellow ferric chloride aqueous solution.
The NMR-spectroscopy data (Section I-7) were an evidence in favor of N-oxidation with subsequent Cope elimination.
When iron(II) sulfate was added into the xylenol orange aqueous solution, the yellow color and absorption maximum on 430 nm were not changed. Further addition of hydrogen peroxide led to violet coloring with maximal absorption at 575 nm. I-6. Oxidation of ferric pyridoxine chelates. The ferric pyridoxine chelates determine the red color of aqueous solution with absorption maximum of visible light at 449 nm. The hydrogen peroxide adding resulted in solution discoloration. The final solution had no absorption maximums in visible region and it looked like pale yellow ferric chloride aqueous solution.
Based on NMR-spectroscopy results (Section I-7) we concluded that the primary alcohol group at pyridine nucleus 4 carbon atom oxidized into carboxyl group as the main reaction path.
When iron(II) sulfate was added into the pyridoxine aqueous solution, the latter remained colorless. Further addition of hydrogen peroxide led to red coloring with maximal absorption at 449 nm. I-7. The NMR-spectroscopy. We used a Bruker Avance III HD NMR-spectrometer for determining the products of benzylpenicillin and two chelate complexes oxidation. The initial structures of organic molecules are presented on There were two quaternary carbon atoms signals 50.04 and 56.01 ppm in 13 C-NMR spectrum (absent in DEPT-135 spectrum). There were no conditions for new methyl groups and quaternary carbon atoms generation during benzylpenicillin oxidation by hydrogen peroxide. Thus, the increasing of their number was determined by surrounding modification that was the stepwise S-oxidation (sulfide → sulfoxide → sulfone): Oxidation of ferric xylenol orange chelates. The NMR-spectroscopy data are available in our previous paper [11] where the N-oxidation of tertiary amine fragments is argued: The scheme of proposed subsequent N-oxide Cope elimination and some rearrangements with possible oligomerization of intermediates is in the article [11]. I-7-c. Oxidation of ferric pyridoxine chelates. The detailed NMR-spectroscopy data are available in our previous paper [12]. Partially, for pyridoxine aqueous solution, two methylene fragments of primary alcohol groups gave 4.60 and 4.79 ppm signals in 1 H-NMR spectrum and 56.63 and 58.02 ppm signals in 13 C-NMR spectrum (negative in DEPT-135 spectrum). For final solution, formed after hydrogen peroxide adding to the red aqueous solution of ferric pyridoxine chelates, there were signals 3.97 ppm in 1 H-NMR spectrum and 59.04 ppm in 13 C-NMR spectrum (negative in DEPT-135 spectrum) that were for only one methylene group. We concluded that the primary alcohol group at para-position relative to pyridine nitrogen atom was oxidized into carboxyl group. The acylation of hydroxyl group at nearby molecule 3 carbon atom resulted in intermolecular ester bond formation. That was the main reaction pathway (Fig. 4). The additional oxidative modifications were the oxidation of another primary alcohol group at pyridine nucleus 5 carbon atom and the N-oxidation of pyridine nitrogen atom [12].
I-8. Electro-Fenton reactions. Earlier, Chumakov A.A. et al. [13] reproduced electro-Fenton-like reactions of metal ions Cr 3+ , Fe 2+ , Fe 3+ , Co 2+ , Ni 2+ , Cu 2+ , La 3+ , and Ce 4+ with electrogenerated hydrogen peroxide. The voltammetry with mercury film working electrode was used. The molecular dioxygen was electrochemically reduced on the mercury film electrode. The voltammogram of oxygen reduction was a two-wave curve (Fig. 5a). The first wave in the voltage ranging from zero to -1 V corresponded to hydrogen peroxide generation reaction [13]: The second wave between -1 and -2 V was the electric current of two-electron hydrogen peroxide reduction: When hydrogen peroxide was added into an electrochemical cell, we watched amperage increasing between -1 and -2 V. The electric current was proportional to hydrogen peroxide concentration in the cell, and there were no changes in the shape of the second wave of voltammetric curve (Fig. 5b). In the presence of Cr 3+ , Fe 2+ , Fe 3+ , Co 2+ , Ni 2+ , Cu 2+ , La 3+ , or Ce 4+ ions, there was appearance of a new wave in the voltage range between -1.4 and -1.8 V with current maximum near -1.6 V (Fig. 5c).

Fig. 5. Voltammograms: a) oxygen reduction on the mercury film electrode, b) changing when hydrogen peroxide is added, c) changing in the presence of transition metal ions
The voltammogram changing in the presence of transition metal ions (Fig. 5c) was interpreted as hydroxyl radical generation and its one-electron cathodic reduction [13]. However, at the present time we outline another explanation presented in the discussion section of this paper.
I-9. Theoretical foundation. Theoretically, hydrogen peroxide molecule can participate not only in one-electron redox reactions (Equations (1) and (2)) but also in proton transfer reactions with generation of hydroperoxide anion [14], hydroperoxonium cation [15] Noncatalytically, the oxywater generation can take place in an associate of two hydrogen peroxide molecules, due to simultaneous intermolecular proton transfers (Fig. 6). The oxywater is a bipolar ion, in which the opposite charges are localized on neighboring oxygen atoms. This determines the displacement of electron density from the negatively charged atom to the positively charged atom. As a result, interoxygen bond heterolytically dissociates (intramolecular disproportionation) with liberation of a water molecule and formation of an oxygen atom (oxene) in a 1 Dsinglet quantum state. This atom has a vacant atomic orbital [3]: II. Discussion II-1. Common basis for oxywater-oxene conception. The results of thermodynamic calculations are the evidence for impossibility of spontaneous hydroxyl and hydroperoxyl radicals generation from hydrogen peroxide in aqueous solution, even in the presence of iron and copper ions, because radicalgenerating reactions are endergonic (Table 2). However, the facts of iron and copper ions interactions with hydrogen peroxide are undoubted. Firstly, there is a great number of the literature data, for example, referenced in papers [1,3]. Secondly, we have shown that hydrogen peroxide spontaneously decomposes in ferrous, ferric, and cupric Fenton reaction systems (Fig. 2). Thirdly, ferric xylenol orange and ferric pyridoxine complexes are oxidized rapidly and spontaneously by hydrogen peroxide, due to interaction of chelated iron(III) ions with oxidant molecules (Sections I-5 and I-6). Lastly, hydrogen peroxide in aqueous solutions spontaneously (without thermal, photonic or catalytic activation) oxidizes the sulfur atoms of hyposulfite and benzylpenicillin molecules (Sections I-3 and I-4). Thus, a hydrogen peroxide molecule generates another intermediate that differs from hydroxyl and hydroperoxyl radicals. In our opinion, the hydrogen peroxide molecule isomerization to oxywater zwitterion (Equation (15)) with subsequent oxywater intramolecular disproportionation (Equation (17)) is a process that is very suitable for explaining all events of hydrogen peroxide reactivity in aqueous systems.
II-2. Hydrogen peroxide decomposition. In contrast to widespread schemes of one-electron redox reactions (Equations (1) and (2)), we postulate the priority of proton transfer reactions in hydrogen peroxide aqueous solutions (Equations (13)- (15)). In contrast to our previous assumption for thermal or photonic one-electron with one-proton transfer in the associate of two hydrogen peroxide molecules with simultaneous hydroxyl and hydroperoxyl radicals generation (Fig. 1), we now argue the oxywater zwitterions generation in hydrogen peroxide dimer, due to simultaneous intermolecular proton transfers (Fig. 6). Further, the oxygen atoms in 1 D-singlet quantum state are generated (Equation (17)). For Fenton reaction systems, the zwitterionization of hydrogen peroxide in Lewis acid-base complex with Me n+ (Equation (16)) is also followed by intramolecular disproportionation of oxywater (Equation (17) . In contrast to our previous argumentation for the free radical chain mechanism of hydrogen peroxide decomposition [9], we now substantiate a 1 D-oxene-mediated pathway of hydrogen peroxide disproportionation in aqueous solutions. A singlet oxygen atom interacts with another hydrogen peroxide molecule through targeting the unshared electron pair of oxygen atom by the vacant atomic orbital: The process may be called O-oxidation of hydrogen peroxide. The proton transfer follows this reaction and results in trioxidane (dihydrogen trioxide) formation: Hydrogen trioxide rapidly decomposes and produces water and singlet dioxygen [17]: Thus, dismutation of hydrogen peroxide (Equation (8) The decomposition of hydrogen peroxide in ferric (Fig. 2b) and cupric (Fig. 2c) systems occurs via Equations (16)- (20).
Previously, Chumakov A.A. et al. [3] carried out the substantiation of the singlet quantum state of molecular oxygen generated during hydrogen peroxide decomposition (Equation (21)), using simple quantum chemical graphical modeling. Besides, we have suggested a mechanism of electron spin rotation during the quenching process 1 O 2 → 3 O 2 . We have assumed the formation of an associate ( 1 O 2 ) 2 from antipodes of orbital parameter. Two simultaneous redox reactions (the electron exchange interaction) result in two triplet dioxygen molecules generation. The first molecule 3 O 2 has +1 total electron spin and the second one has -1 total electron spin: The overall equation of hydrogen peroxide disproportionation is the following:

II-3. Classic Fenton reaction mechanism.
The trivalent iron and divalent copper ions serve as catalysts of oxywater formation and singlet 1 D-oxygen atom generation. The ferric ions are more active than cupric ions (Fig. 2d). The ferric and cupric ions are true catalysts because their oxidation states do not change after the completion of hydrogen peroxide decomposition. In contrast, ferrous iron ions are not true catalysts of hydrogen peroxide dismutation because of change of their oxidation state to ferric. Ferrous sulfate is a reagent oxidized by hydrogen peroxide: The process of hydrogen peroxide decomposition in ferrous Fenton system (Fig. 2a) is accompanied by precipitation of iron(III) oxide-hydroxide with rusty color:  2 3   (26) The oxidation of divalent iron to trivalent one by hydrogen peroxide is proved by violet coloration (λ max =575 nm) of yellow ferrous xylenol orange aqueous solution (Section I-5) and red coloration (λ max =449 nm) of colorless ferrous pyridoxine aqueous solution (Section I-6), when hydrogen peroxide is added.
We argued [3] the rapid and inevitable one-electron transfer within the iron(II)-oxene complex: The ferric-oxyl radical-anion complex is known to be generated from nitrous oxide on the surface of FeZSM-5 zeolite and is considered as α-oxygen complex [18]. In our opinion, the classic Fenton reaction occurs through this complex formation: Such view is alternative to the widespread classic Fenton reaction conceptions of hydroxyl-free radical generation (Equation (3)) or the oxoferryl(IV) cation formation by Equation (29) The α-oxygen complex oxidizes the second ferrous ion: Further, the hydrolysis of complex cation occurs:  (31) is the mechanism of reaction by Equation (24). The generated ferric ions decompose hydrogen peroxide by the mechanism including consequent formation of oxywater, 1 D-oxene, dihydrogen trioxide, singlet dioxygen and quenching of singlet dioxygen to triplet quantum state (Equations (16)- (20), and (22)). The rate of hydrogen peroxide decomposition in ferrous system decreases compared to ferric system, because of removal of Fe 3+ ions in the first one by precipitation of iron(III) oxide-hydroxide (Fig. 2).

II-4. Electro-Fenton reactions.
Earlier [13], we interpreted changes of the second voltammogram wave (Fig. 5c) as hydroxyl radical generation and its one-electron cathodic reduction including redox cycling of metal ion oxidation state: (34) At the present time, we maintain another explanation. The second wave is the electric current of two-electron hydrogen peroxide reduction (Equation (12)), but the detailed mechanism includes hydrogen peroxide isomerization to oxywater (Equation (15)) and 1 D-oxene generation (Equation (17)). The polarization of hydrogen peroxide occurs in an electrochemical cell under the electric field between electrodes. Thus, the second wave is in truth the current of 1 D-oxene two-electron reduction:  (Fig. 5c) is the current of oxyl radical-anion one-electron cathodic reduction (versus Equation (33)): The hydrolysis of oxide anion occurs: The redox cycling of metal ion oxidation state occurs as Me n+ ↔Me n+1 . For the used salts [13], there are cycles Fe 2+ ↔Fe 3+ , Cu + ↔Cu 2+ , La 2+ ↔La 3+ , Ce 3+ ↔Ce 4+ , and others.

II-5. Model substrates oxidation.
The sulfur atom of benzylpenicillin molecule has two unshared electron pairs. Thus, the S-oxidation by Equation (9) occurs via targeting of sulfur atom electron pairs by vacant atomic orbitals of two singlet oxygen atoms generated from hydrogen peroxide molecules. The reaction of hyposulfite oxidation by hydrogen peroxide occurs rapidly and spontaneously. As supported by us, the reaction occurs via the non-radical stepwise S-oxidation by 1 D-oxene oxidant and results in dithionite, metabisulfite, dithionate, pyrosulfate, and bisulfate formation: For ferric chelate oxidation, the direct oxidant is also a singlet 1 D-oxygen atom generated by the chelated Fe 3+ ions. The N-oxidation of xylenol orange molecule (Equation (10)) occurs via targeting of the nitrogen atom unshared electron pair by 1 D-oxene vacant atomic orbital [11]. The oxidation of primary alcohol group of pyridoxine (Fig. 4) There is consequent formation of geminal diol, aldehyde, and carboxylic acid. Further, the acylation of hydroxyl group at nearby molecule 3 carbon atom results in intermolecular ester bond formation [12].
II-6. The application of oxywater-oxene conception. The oxywater-oxene concept is successfully applicable to explain the catalytic activity of redox-inactive substances, for which the free radical or high-valence species generation schemes are unusable. For instance, gallium(III) and aluminum(III) nitrates catalyze the epoxidation of some olefins with hydrogen peroxide [19]. The mechanism is unclear. We assume the formation of Lewis acid-base complexes [Al 3+ O 0 ( 1 D)] 3+ and [Ga 3+ O 0 ( 1 D)] 3+ . The epoxidation occurs by singlet oxene. For another example, the hydrogen peroxide lanthanum(III) system (including La(NO 3 ) 3 , or La(OH) 3 , or La 2 O 3 ) is a generator of 1 ∆ g -singlet dioxygen [20]. Although we maintain the redox cycling La 2+ ↔La 3+ in an electrochemical cell [13], it is unlikely that lanthanum(III) changes oxidation degree +3 when generating a singlet dioxygen 1 O 2 from hydrogen peroxide [20]. We suggest the formation of a similar Lewis acid-base complex [La 3+ O 0 ( 1 D)] 3+ and subsequent generation and decomposition of dihydrogen trioxide (Equations (18)- (20)).
Hydroperoxides, including the simplest hydrogen peroxide, are widely used in organic synthesis for oxygen functionalization processes, such as alkanes and arenes hydroxylation, alkenes epoxidation, Baeyer-Villiger ketones oxidation to esters, organonitrogen compounds N-oxidation and organosulfur compounds S-oxidation [3]. In addition to monooxygen oxidation, the dioxygen functionalization by 1 ∆ g -singlet molecular oxygen also takes place in organic synthesis and includes synthesis of hydroperoxides from alkenes and cyclic peroxides from alkadienes [21][22][23].
We used our oxywater-oxene concept for explanation of the mechanisms of hydroperoxide monooxygen (Fig. 7) and dioxygen (Fig. 8) Alkenes epoxidation via targeting of pi-electron pair by vacant atomic orbital of 1 D-oxene Baeyer-Villiger ketone oxidation to ester via insertion of oxygen atom between carbonyl and alkyl groups as result of targeting of sigma-electron pair by vacant atomic orbital of 1 D-oxene. The carbon atom of carbonyl group has a partial positive charge, which coordinates an oxene The N-and S-oxidation of organonitrogen and organosulfur compounds via targeting of unshared electron pairs of heteroatoms by vacant atomic orbital of 1

Conclusion
Hydrogen peroxide isomerization to oxywater zwitterion with subsequent oxywater intramolecular disproportionation is a process that is very suitable for explaining all events of hydrogen peroxide decomposition and oxidative reactivity in aqueous systems.
An oxygen atom (oxene) in a singlet quantum state has a vacant atomic orbital. It mediates disproportionation of hydrogen peroxide in aqueous solutions via O-oxidation of the second hydrogen peroxide molecule. The formed trioxidane (dihydrogen trioxide) rapidly decomposes and produces water and singlet dioxygen. The quenching of the latter substance into triplet quantum state occurs via electron exchange interaction between two singlet dioxygen molecules.