Hubbry Logo
TrioxidaneTrioxidaneMain
Open search
Trioxidane
Community hub
Trioxidane
logo
8 pages, 0 posts
0 subscribers
Be the first to start a discussion here.
Be the first to start a discussion here.
Trioxidane
Trioxidane
Trioxidane
View on Wikipedia
from Wikipedia
Trioxidane
Structural formula of trioxidane
Structural formula of trioxidane
Names
Preferred IUPAC name
Trioxidane (only preselected name)[1]
Systematic IUPAC name
Dihydrogen trioxide
Other names
Hydrogen trioxide
Dihydroxy ether
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
200290
  • InChI=1S/H2O3/c1-3-2/h1-2H checkY
    Key: JSPLKZUTYZBBKA-UHFFFAOYSA-N checkY
  • InChI=1/H2O3/c1-3-2/h1-2H
    Key: JSPLKZUTYZBBKA-UHFFFAOYAV
  • OOO
Properties
H2O3
Molar mass 50.013 g·mol−1
Related compounds
Related compounds
 Hydrogen peroxide; Hydrogen ozonide; Hydroperoxyl
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Trioxidane (systematically named dihydrogen trioxide,[2][3]), also called hydrogen trioxide[4][5] is an inorganic compound with the chemical formula H[O]
3H (can be written as [H(μ-O
3)H] or [H
2O
3]). It is one of the unstable hydrogen polyoxides.[4] In aqueous solutions, trioxidane decomposes to form water and singlet oxygen:

Reaction of trioxidane (blue) with water (red) results in decomposition to oxygen and an additional water molecule.

The reverse reaction, the addition of singlet oxygen to water, typically does not occur in part due to the scarcity of singlet oxygen. In biological systems, however, ozone is known to be generated from singlet oxygen, and the presumed mechanism is an antibody-catalyzed production of trioxidane from singlet oxygen.[2]

Preparation

[edit]

Trioxidane can be obtained in small, but detectable, amounts in reactions of ozone and hydrogen peroxide, or by the electrolysis of water. Larger quantities have been prepared by the reaction of ozone with organic reducing agents at low temperatures in a variety of organic solvents, such as the anthraquinone process. It is also formed during the decomposition of organic hydrotrioxides (ROOOH).[3] Alternatively, trioxidane can be prepared by reduction of ozone with 1,2-diphenylhydrazine at low temperature. Using a resin-bound version of the latter, relatively pure trioxidane can be isolated as a solution in organic solvent. Preparation of high purity solutions is possible using the methyltrioxorhenium(VII) catalyst.[5] In acetone-d6 at −20 °C, the characteristic 1H NMR signal of trioxidane could be observed at a chemical shift of 13.1 ppm.[3] Solutions of hydrogen trioxide in diethyl ether can be safely stored at −20 °C for as long as a week.[5]

The reaction of ozone with hydrogen peroxide is known as the "peroxone process". This mixture has been used for some time for treating groundwater contaminated with organic compounds. The reaction produces H2O3 and H2O5.[6]

Structure

[edit]

In 1970–75, Giguère et al. observed infrared and Raman spectra of dilute aqueous solutions of trioxidane.[4] In 2005, trioxidane was observed experimentally by microwave spectroscopy in a supersonic jet. The molecule exists in a skewed structure, with an oxygen–oxygen–oxygen–hydrogen dihedral angle of 81.8°. The oxygen–oxygen bond lengths of 142.8 picometer are slightly shorter than the 146.4 pm oxygen–oxygen bonds in hydrogen peroxide.[7] Various dimeric and trimeric forms also seem to exist.

There is a trend of increasing gas-phase acidity and corresponding pKa as the number of oxygen atoms in the chain increases in HOnH structures (n=1,2,3).[8]

Reactions

[edit]

Trioxidane readily decomposes into water and singlet oxygen, with a half-life of about 16 minutes in organic solvents at room temperature, but only milliseconds in water. It reacts with organic sulfides to form sulfoxides, but little else is known of its reactivity.

Recent research found that trioxidane is the active ingredient responsible for the antimicrobial properties of the well known ozone/hydrogen peroxide mix. Because these two compounds are present in biological systems as well it is argued that an antibody in the human body can generate trioxidane as a powerful oxidant against invading bacteria.[2][9] The source of the compound in biological systems is the reaction between singlet oxygen and water (which proceeds in either direction, of course, according to concentrations), with the singlet oxygen being produced by immune cells.[3][10]

Computational chemistry predicts that more oxygen chain molecules or hydrogen polyoxides exist and that even indefinitely long oxygen chains can exist in a low-temperature gas. With this spectroscopic evidence a search for these types of molecules can start in interstellar space.[7] A 2022 publication suggested the possibility of the presence of detectable concentrations of polyoxides in the atmosphere.[11]

See also

[edit]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Trioxidane

View on Grokipedia
from Grokipedia
Trioxidane is an unstable with the molecular formula H₂O₃, also known as hydrogen trioxide or dihydrogen trioxide. It represents the simplest , structurally analogous to (H₂O₂) but featuring an extended oxygen chain. The adopts a zigzag skew-chain conformation with C₂ , where the O-O-O chain has bond lengths of approximately 1.388 for the terminal O-O bonds and 1.509 for the central O-O bond, as determined by high-level quantum chemical calculations. Its at 298.15 K is -90.27 ± 0.70 kJ/mol, reflecting its thermodynamic instability relative to and oxygen. Trioxidane is more acidic than , with a pKa of 9.5 ± 0.2, and it exhibits hydrogen-bonding capabilities that influence its behavior. Synthesis of trioxidane typically involves low-temperature reactions, such as the ozonation of 1,2-diphenylhydrazine in deuterated acetone at -78 °C, yielding concentrations up to 0.1 M, or the peroxone process combining and . It decomposes via first-order kinetics primarily to and singlet oxygen (¹O₂), with an activation enthalpy of 17.7–18.8 kcal/mol; the is 16 ± 2 minutes in acetone-d₆ at 20 °C but only ~20 ms in , where decomposition is catalyzed by or . Stability is enhanced in organic solvents or oxygen-containing bases, and it forms hydrogen-bonded complexes that mitigate rapid breakdown. Trioxidane plays a role in atmospheric and environmental oxidation processes, including the degradation of pollutants and reactions, and has been implicated as an intermediate in biological systems such as antibody-catalyzed water oxidation to produce and . Recent advances highlight its formation in extraterrestrial ice analogs under galactic at 5 K, identified via with an ionization energy of 10.95 ± 0.05 eV, expanding its relevance to prebiotic chemistry on icy bodies like Europa and comets. Comprehensive reviews underscore ongoing progress in its spectroscopic characterization, theoretical modeling, and potential applications in selective oxidation despite persistent challenges with isolation.

General Properties

Definition and Nomenclature

Trioxidane is an with the H₂O₃, often represented in alternative notations as H[O]₃H or HOOOH to emphasize its linear oxygen chain structure. Its is 50.014 g·mol⁻¹. The for this compound is trioxidane, while its systematic IUPAC name is μ-trioxidanediidodihydrogen, reflecting the bridged peroxide-like linkages between the oxygen atoms and substituents. The term "trioxidane" derives from the Greek prefix "tri-" indicating three oxygen atoms in the chain, analogous to naming conventions for polyoxides. Trioxidane belongs to the class of , a series of unstable compounds extending beyond (H₂O₂) and (H₂O), characterized by extended chains of oxygen atoms bonded to . Unlike the more stable H₂O₂, trioxidane exhibits significant instability, with a short under typical conditions.

Physical Characteristics

Trioxidane has not been isolated as a pure substance due to its extreme instability, but it manifests as a transient in low-temperature matrices or solutions, where it appears as a colorless material akin to related polyoxides like . Owing to rapid decomposition, empirical measurements of and boiling points are unavailable; trioxidane begins to decompose at approximately -55 °C when prepared as a glassy solid, precluding observation of a distinct transition, while at it persists only briefly in suitable media. The of pure trioxidane is theoretically estimated at 1.65 g/cm³ at -20 °C, derived from molecular volume calculations. Trioxidane exhibits high in both and organic solvents, including acetone and , reflecting its polar nature; however, it decomposes with a of about 20 ms in aqueous environments, contrasting with longer stability ( of 16 minutes at 20 °C) in deuterated acetone. Thermodynamic data indicate a (Δ_f H°) of -90.60 ± 0.40 kJ/mol for gaseous trans-trioxidane at 298.15 , as determined by the Active Thermochemical Tables (ATcT).

Stability and Detection

Trioxidane, or dihydrogen trioxide (HOOOH), exhibits significant instability, with its varying markedly depending on the environment. In organic solvents such as acetone-d₆, it has a half-life of 16 ± 2 minutes at 20 °C, while in deuterated acetone (CD₃COCD₃) at -20 °C, the half-life extends similarly under controlled conditions. In aqueous solutions, however, occurs rapidly, with a half-life of approximately 20 milliseconds at . Stability is influenced by , polarity, and ; lower temperatures and nonpolar organic solvents enhance persistence, whereas polar protic solvents like accelerate breakdown, and optimal stability in aqueous media occurs around 1.7 (0.02 M H⁺), yielding a half-life of up to 2 seconds. The primary decomposition pathway of trioxidane yields water and (¹Δ_g O₂), which is a reactive of molecular oxygen. The dominant process is the unimolecular elimination to H₂O and ¹O₂, reflecting the weak central O–O bond. Detection of trioxidane relies on techniques due to its transient nature. (¹H NMR) identifies it via a characteristic signal at 13.1 ppm in acetone-d₆ at -20 °C, downfield from typical OH protons due to the adjacent linkages. (IR) reveals the antisymmetric O–O stretch at 776 cm⁻¹, a marker band distinguishable even in aqueous matrices. , pioneered in 1970s studies, confirms skeletal vibrations, including O–O stretches around 800–900 cm⁻¹ in matrix-isolated samples, providing complementary evidence of its structure. Due to its short across conditions, trioxidane cannot be isolated or stored and must be generated in situ for experimental studies, often via low-temperature ozonation of or related precursors. This limitation underscores the challenges in handling it, restricting applications to transient reactive intermediates.

Synthesis

Laboratory Preparation

Trioxidane (H₂O₃), also known as dihydrogen trioxide, is primarily prepared in settings through controlled reactions that generate it in trace to millimolar quantities for spectroscopic analysis, often under cryogenic conditions to enhance stability. The most common method involves the reaction of ozone (O₃) with hydrogen peroxide (H₂O₂), known as the Peroxone process, which proceeds via the equation H₂O₂ + O₃ → H₂O₃ + O₂. This thermal reaction is typically conducted in deuterated solvents such as acetone-d₆ or tetrahydrofuran-d₁₀ at -78 °C under an inert atmosphere to minimize decomposition. Yields reach approximately 2.9 mM of H₂O₃, detectable by ¹H NMR spectroscopy at δ = 13.6 ± 0.2 ppm. An alternative low-temperature approach utilizes the ozonation of organic reducing agents, such as 1,2-diphenylhydrazine, in solvents like acetone-d₆ at -78 °C. This generates H₂O₃ alongside byproducts including 1,2-diphenyldiazene and hydrogen peroxide, with concentrations up to 0.1 M achievable and confirmed via ¹H, ²H, and ¹⁷O NMR spectroscopy. The process requires an inert atmosphere and rapid cooling. Trace amounts of trioxidane can also be produced via the electrical discharge dissociation of , followed by chilling to -190 °C to form a glassy solid containing less than 5% H₂O₃. Detection is performed using , and the method necessitates an inert environment to isolate the unstable product. These preparations generally yield low concentrations at the ppm to low mM level, demanding immediate spectroscopic characterization due to the compound's instability.

Industrial and Alternative Methods

Trioxidane is generated continuously in the peroxone process, which combines and for advanced oxidation in applications, such as the remediation of contaminated and . In this method, trioxidane forms as an intermediate through the reaction of with hydroperoxide ions, contributing to the production of hydroxyl radicals without requiring isolation due to its transient nature. The process operates under ambient conditions in flow systems, enabling scalable treatment of pollutants like volatile organic compounds. Alternative production routes emphasize in situ generation to leverage trioxidane's reactivity while mitigating . Recent electrochemical approaches pair the simultaneous production of and in specialized reactors, enhancing peroxone efficiency for continuous oxidation processes. These methods avoid direct handling of trioxidane, focusing instead on controlled delivery of precursors for on-demand formation. Industrial-scale isolation of trioxidane remains unfeasible owing to its , with half-lives on the order of milliseconds in aqueous environments and up to 16 minutes in organic solvents at . Consequently, it is employed exclusively for oxidative applications, such as , where its role in radical generation provides environmental benefits without purification steps. Innovations in trioxidane handling include cryogenic trapping at -78°C during ozonation reactions, achieving concentrations up to 0.1 M in solvents like acetone-d6 for spectroscopic study, though not for commercial production. Matrix isolation techniques at even lower temperatures have further enabled detailed examination of its structure and reactivity, supporting fundamental research into higher .

Molecular Structure

Geometry and Bonding

Trioxidane adopts a skewed, linear geometry described as H-O-O-O-H, featuring C₂ in its most stable trans conformation. This arrangement results in two equivalent but oppositely oriented O-O-O-H dihedral angles of 81.3° each, distinguishing it from the more planar 112° dihedral in . The structure is characterized as a skew-, where between oxygen lone pairs and the σ* orbitals of O-H bonds stabilizes the molecule and influences its overall shape. The O-O bond lengths in trioxidane differ between terminal and central bonds: the terminal O-O bonds are approximately 138.8 pm, while the central O-O bond is about 150.9 pm, with the terminal bonds shorter than the 145.8 pm O-O bond in , reflecting enhanced bonding due to the hyperconjugative effects in the extended chain. In contrast, the terminal O-H bonds remain typical for peroxides, approximately 96 pm, with no significant deviation reported in high-level calculations. The bonding in the O-O linkages exhibits weak, peroxide-like single-bond character, as evidenced by (DFT) and coupled-cluster methods like CCSD(T), which predict bond orders close to 1 but with elongated distances indicative of low dissociation energies around 20-30 kcal/mol. Among isomers, the trans configuration is strongly preferred, lying 9 kJ/mol below the cis form (Cₛ ), with a low barrier of 22 kJ/mol for interconversion; no stable cyclic exists, as the [O(H₂O)O] is destabilized by 485 kJ/mol relative to the trans . Quantum chemical predictions confirm the trans skewed as the global energy minimum, with CCSD(T)/aug-cc-pVTZ calculations yielding vibrational frequencies that match experimental data, including an antisymmetric O-O stretch at 776 cm⁻¹. These models, employing basis sets like aug-cc-pVTZ and complete basis set extrapolations, underscore the partial single-bond nature of the O-O interactions through analysis of and natural bond orbitals.

Spectroscopic Evidence

Nuclear magnetic resonance (NMR) has provided key evidence for the structure of trioxidane (HOOOH). The ¹H NMR spectrum exhibits a characteristic downfield singlet at 13.1 ppm in acetone-d₆ at -20°C, attributed to the -OOH protons, which experience deshielding due to the adjacent oxygen atoms. This signal shifts slightly to 13.6 ± 0.2 ppm at -78°C under similar conditions. Additionally, ¹⁷O NMR confirms the presence of distinct oxygen environments, with chemical shifts assigned to the terminal and central oxygens, matching calculated values from GIAO/CCSD(T)/qz2p methods and supporting the H-O-O-O-H connectivity. Infrared (IR) and have been instrumental in identifying trioxidane through its vibrational signatures, particularly the O-O stretching modes around 800 cm⁻¹. Early studies in the by Giguère et al., involving the reaction of with in frozen water-peroxide matrices at temperatures, revealed IR absorptions assigned to O-O vibrations in H₂O₃; for the deuterated isotopomer, bands at 857, 820, and 760 cm⁻¹ shifted to 806, 775, and 717 cm⁻¹ upon ¹⁸O substitution, confirming the assignments. More recent argon matrix isolation IR spectra identified the antisymmetric O-O stretch at 776 cm⁻¹ for the normal isotopomer, consistent with ab initio predictions and further validating the skewed structure. of matrix-isolated or glassy solid samples similarly show fundamental skeletal vibrations in this region, providing complementary evidence for the peroxide linkages. Microwave spectroscopy offered the first gas-phase structural confirmation of trioxidane in 2005, using techniques on samples generated in a supersonic jet via pulsed discharge. Rotational constants derived from the spectra of HOOOH and its DOOOD isotopomer precisely match those expected for a skewed analyzed as having C₁ , consistent with a trans-like O-O-O backbone theoretically predicted to have C₂ , enabling accurate predictions for atmospheric . Mass spectrometry under low-pressure conditions has detected the parent ion of trioxidane at m/z 50, corresponding to H₂O₃⁺, with an ionization onset at 10.95 ± 0.05 eV in photoionization experiments on ice analogs, supporting its identification alongside isotopic variants.

Chemical Reactivity

Decomposition Pathways

Trioxidane, or dihydrogen trioxide (HOOOH), primarily undergoes unimolecular self-decomposition to form water and singlet oxygen according to the reaction: HOOOH→H2O+1O2\mathrm{HOOOH \rightarrow H_2O + ^1O_2} This pathway is supported by experimental evidence from NMR spectroscopy and product analysis in organic solvents. The decomposition follows first-order kinetics, with a half-life of approximately 16 minutes at 20°C in solvents such as acetone-d₆, methyl acetate, and tert-butyl methyl ether. In aqueous environments, the half-life is significantly shorter, around 20 milliseconds, due to catalytic effects. The temperature dependence of the decomposition rate indicates an activation energy of 15–19 kcal/mol, varying slightly with solvent and catalytic conditions; for water-assisted pathways, values range from 17.3 to 18.8 kcal/mol. This energy barrier aligns with the instability of the weak O–O bonds in the trioxidane structure. The mechanism involves two successive water-assisted homolytic cleavages of the O–O bonds, leading to the formation of singlet oxygen without significant radical chain propagation under typical conditions. At higher temperatures or in the gas phase, homolytic O–O cleavage can generate hydroxyl (•OH) and hydroperoxyl (•OOH) radicals, potentially initiating minor radical pathways. Decomposition is accelerated by protic catalysts; water and act as bifunctional catalysts, lowering the through hydrogen bonding and proton transfer. Acidic conditions further enhance the rate via of trioxidane to form HOOOH₂⁺, facilitating O–O bond scission. In contrast, oxygen-containing bases like ethers stabilize trioxidane by forming hydrogen-bonded complexes that inhibit the . The production of in this process contributes to its potential role in oxidative applications.

Reactions with Other Substances

Trioxidane, also known as dihydrogen trioxide (H₂O₃), demonstrates selective oxygen transfer in oxidation reactions with organic sulfides, converting them to sulfoxides via a mild electrophilic mechanism. For instance, substituted phenyl methyl sulfides are oxidized to their corresponding sulfoxides in solvents like acetone or at -40°C, with a Hammett showing ρ = -1.90 ± 0.02, indicative of sulfur's nucleophilic attack on the central oxygen atom of H₂O₃. The general reaction can be represented as: H2O3+R2S→R2SO+H2O2\mathrm{H_2O_3 + R_2S \rightarrow R_2SO + H_2O_2} This process contrasts with hydrogen peroxide, which does not react under identical conditions, underscoring trioxidane's superior reactivity as an oxidant. In the Peroxone process, generated from the reaction of ozone with hydrogen peroxide, trioxidane serves as a key intermediate in a chain mechanism that enhances hydroxyl radical (•OH) production for pollutant degradation. The formation of H₂O₃ from the initial O₃–H₂O₂ complex facilitates subsequent radical propagation, where its decomposition contributes to amplifying •OH yields beyond direct O₃ consumption. This role positions trioxidane as an electrophilic oxidant (with nucleophilicity parameter X_Nu = 0.17 ± 0.01 toward thianthrene 5-oxide), ultimately reducing to H₂O₂ or H₂O in redox cycles.

Applications and Biological Role

Environmental and Industrial Uses

Trioxidane plays a key role in the peroxone process, an advanced oxidation method where and react in aqueous solutions to generate reactive species for . In this process, trioxidane forms as an intermediate through the interaction of with , contributing to the degradation of contaminants such as volatile organic compounds (VOCs) in . The peroxone approach enhances production, enabling the breakdown of persistent pollutants like and that are resistant to individual or treatments alone. The efficiency of the peroxone process surpasses that of ozonation or oxidation independently, with yields reaching approximately 50% under optimal conditions, allowing for faster and more complete contaminant mineralization. This has led to its adoption in municipal water systems since the 1980s, particularly for treating sources contaminated with organic pollutants. In practice, the process is implemented by injecting into ozonated , optimizing radical generation while minimizing residual . Ozone-hydrogen peroxide systems are used for oxidation during bleaching of thermomechanical pulp. This combined treatment improves pulp brightness by up to 4.3% in sequences involving alkaline extraction, offering a chlorine-free alternative that reduces effluent compared to traditional methods. The system's selectivity is enhanced by additives like , which minimize fiber yield loss while promoting sustainable delignification. Trioxidane's environmental impact is favorable, as it decomposes primarily into and without forming harmful byproducts, supporting its use in disinfection where provides additional activity against and viruses. The peroxone process has been integrated into ozonation facilities for effective removal, aiding compliance with stringent EU standards.

Biomedical Implications

Trioxidane, or dihydrogen trioxide (H₂O₃), plays a role in immune responses through antibody-catalyzed formation from (H₂O₂) and molecular oxygen (O₂), or more precisely via the reaction of (¹O₂) with . This process, known as the antibody-catalyzed water oxidation pathway, generates H₂O₃ as a transient intermediate that decomposes to H₂O₂, enhancing action against pathogens. All immunoglobulins, irrespective of antigenic specificity, exhibit this catalytic activity, suggesting a universal mechanism in host defense. In neutrophils, surface-bound antibodies facilitate H₂O₃ production during the oxidative burst, where ¹O₂ generated by NAD(P)H oxidase reacts to form H₂O₃, which decomposes to water and , contributing to the overall release of ROS including and . This amplifies , enabling efficient killing by disrupting microbial membranes and promoting via induction, such as TNF-α and IL-8. However, excessive H₂O₃-derived oxidants may exacerbate tissue damage in chronic conditions like or . Low concentrations of H₂O₃ are non-toxic and support cellular signaling, but elevated levels induce cytotoxicity via and protein oxidation, contributing to pathological oxidative damage. Research from 2005 highlights H₂O₃'s stability and detection in biological media, such as organic solvents mimicking cellular environments, confirming its role as an oxidant intermediate with implications for both protective immunity and disease pathology like cancer. Spectroscopic evidence (NMR, IR) supports its transient presence ( ~16 min in acetone-d₆ at 20°C), underscoring the need for precise control in biological contexts.

Historical Development

Early Observations

The study of emerged as an extension of post-World War II research on peroxides, driven by applications in propulsion systems, , and oxidative processes, where high-concentration served as a key precursor and model compound. This broader interest in oxygen-rich species prompted investigations into higher polyoxides like trioxidane (H₂O₃), initially postulated as transient intermediates in reactions involving and or . The first spectroscopic evidence for trioxidane came from infrared (IR) spectroscopy in 1970, when Paul A. Giguère and Kazimiera Herman identified matrix-stabilized intermediates in the products of electrical discharges through water vapor at low temperatures (around -196 °C). These experiments revealed characteristic O-O stretching vibrations consistent with H₂O₃ and the related H₂O₄, formed via condensation of peroxy radicals in amorphous ice matrices. Subsequent Raman spectroscopy in 1971 by the same group confirmed these assignments through isotopic labeling with ¹⁸O, observing bands at 500 cm⁻¹ and 760 cm⁻¹ attributable to skeletal vibrations of H₂O₃, further supporting its presence in oxygen-rich H₂O-O₂ mixtures trapped at -180 °C. Between 1970 and 1975, additional IR and Raman studies extended these findings to dilute aqueous solutions generated from ozone-hydrogen peroxide reactions, providing the earliest direct spectral signatures of the molecule. In the 1970s literature, the compound was consistently referred to as "hydrogen trioxide" to denote its structure as H-O-O-O-H, distinguishing it from simpler peroxides and emphasizing its polyoxide nature. This work underscored trioxidane's role in decomposition pathways, though its extreme instability—decomposing rapidly above -100 °C—limited observations to cryogenic conditions. Early characterization efforts were hampered by trioxidane's short lifetime (milliseconds at best in solution) and tendency to disproportionate into and oxygen, necessitating indirect detection via vibrational spectroscopy in matrices or transient species. No bulk isolation was achieved during this period; the molecule remained confined to spectroscopic identification until matrix isolation techniques in the enabled more stable trapping and structural confirmation. These challenges reinforced trioxidane's status as a reactive intermediate rather than a entity, guiding subsequent polyoxide research toward theoretical modeling and advanced trapping methods.

Modern Advances

In the early 2010s, significant progress was made in the synthesis of trioxidane (H₂O₃), enabling its preparation in larger quantities for detailed study. A notable method involved the discharge dissociation of O₂/H₂ gas mixtures at low (~1 ), producing peroxy radical condensates containing up to 30 mol% H₂O₃ and 20 mol% H₂O₄, alongside H₂O₂ and H₂O. These condensates were characterized by , revealing key vibrational modes for H₂O₃ at 500 cm⁻¹ (OOO bend), 756 cm⁻¹ (asymmetric OO stretch), and 878 cm⁻¹ (symmetric OO stretch), confirming its structural integrity in low-temperature solids. A comprehensive review in 2013 highlighted advances in trioxidane's formation via ozonation of hydrazines, such as the low-temperature (–78 °C) reaction with 1,2-diphenylhydrazine yielding up to 0.1 M solutions, and matrix isolation techniques, which allowed infrared spectroscopic detection and theoretical modeling of its C₂-symmetric structure. Reactivity studies emphasized its role as an intermediate in atmospheric oxidation processes and peroxy radical chemistry, with computational methods like CCSD(T) providing insights into its metastable nature and decomposition barriers around 20-25 kcal/mol. These developments built on earlier work, shifting focus from elusive detection to controlled generation and mechanistic understanding. More recent investigations have explored trioxidane's environmental and applied roles. In 2021, research demonstrated its generation via photolysis of H₂O₂/O₃ mixtures under , enhancing the sanitization efficacy of vaporous against bacterial spores like , achieving complete inactivation at reduced H₂O₂ concentrations (0.03%) while preserving filter material integrity. This synergistic process attributes improved biocidal performance to H₂O₃ as an active oxidant species. In astrophysical contexts, a 2025 study simulated galactic irradiation of H₂O-O₂ analogs at 5 K, detecting H₂O₃ via synchrotron vacuum ultraviolet with an of 10.95 ± 0.05 eV, underscoring its potential as a prebiotic oxidizer on interstellar grains and icy bodies.

References

  1. https://pubchem.ncbi.nlm.nih.gov/compound/Trioxidane
  2. https://acta-arhiv.chem-soc.si/52/52-1-1.pdf
  3. https://atct.anl.gov/Thermochemical%20Data/version%201.122/species/?species_number=225
  4. https://www.science.org/doi/10.1126/sciadv.adw5720
  5. https://doi.org/10.1021/cr300512s
  6. https://www.chemspider.com/Chemical-Structure.145859.html
  7. https://iupac.qmul.ac.uk/BlueBook/P6a.html
  8. https://dbpedia.org/page/Trioxidane
  9. https://www.russchemrev.org/RCR4744pdf
  10. https://www.webqc.org/compound-H2O3-H2O3.html
  11. https://atct.anl.gov/Thermochemical%20Data/version%201.176/species/?species_number=3084
  12. https://cdnsciencepub.com/doi/10.1139/v70-581
  13. https://www.pnas.org/doi/10.1073/pnas.202596799
  14. https://www.sciencedirect.com/science/article/abs/pii/S0013468625004128
  15. https://doi.org/10.1021/ja0556530
  16. https://cpeo.org/techtree/ttdescript/peroxz.htm
  17. https://pubs.acs.org/doi/10.1021/es100277d
  18. https://link.springer.com/article/10.1007/s40726-015-0015-z
  19. https://bioresources.cnr.ncsu.edu/resources/improving-bleachability-of-thermomechanical-pulp-by-ozone-hydrogen-peroxide-system/
  20. https://www.lenntech.com/processes/disinfection/chemical/disinfectants-peroxone.htm
  21. https://www.pnas.org/doi/10.1073/pnas.0307311101
  22. https://www.pnas.org/doi/10.1073/pnas.0530251100
  23. https://www.pnas.org/doi/10.1073/pnas.0730791100
  24. https://pmc.ncbi.nlm.nih.gov/articles/PMC3312702/
  25. https://www.sciencedirect.com/topics/earth-and-planetary-sciences/hydrogen-peroxide
  26. https://cdnsciencepub.com/doi/10.1139/v71-364
  27. https://pubs.aip.org/aip/jcp/article/60/1/270/214396/Vibrational-analysis-and-molecular-structure-of
  28. https://chemistry-europe.onlinelibrary.wiley.com/doi/abs/10.1002/ejic.201100767
  29. https://pubs.acs.org/doi/10.1021/cr300512s
  30. https://pubs.rsc.org/en/content/articlehtml/2021/ra/d1ra03675k
Add your contribution
Related Hubs
User Avatar
No comments yet.