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Hydroperoxide
Hydroperoxide
Hydroperoxide
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The general structure of an organic hydroperoxide with the blue marked functional group, where R stands for any group, typically organic

Hydroperoxides or peroxols are compounds of the form ROOH, where R stands for any group, typically organic, which contain the hydroperoxy functional group (−OOH). Hydroperoxide also refers to the hydroperoxide anion (−OOH) and its salts, and the neutral hydroperoxyl radical (•OOH) consist of an unbound hydroperoxy group. When R is organic, the compounds are called organic hydroperoxides. Such compounds are a subset of organic peroxides, which have the formula ROOR. Organic hydroperoxides can either intentionally or unintentionally initiate explosive polymerisation in materials with saturated chemical bonds.[1]

Properties

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The O−O bond length in peroxides is about 1.45 Å, and the R−O−O angles (R = H, C) are about 110° (water-like). Characteristically, the C−O−O−H dihedral angles are about 120°. The O−O bond is relatively weak, with a bond dissociation energy of 45–50 kcal/mol (190–210 kJ/mol), less than half the strengths of C−C, C−H, and C−O bonds.[2][3]

Hydroperoxides are typically more volatile than the corresponding alcohols:

Miscellaneous reactions

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Hydroperoxides are mildly acidic. The range is indicated by 11.5 for CH3OOH to 13.1 for Ph3COOH.[4]

Hydroperoxides can be reduced to alcohols with lithium aluminium hydride, as described in this idealized equation:

4 ROOH + LiAlH4 → LiAlO2 + 2 H2O + 4 ROH

This reaction is the basis of methods for analysis of organic peroxides.[5] Another way to evaluate the content of peracids and peroxides is the volumetric titration with alkoxides such as sodium ethoxide.[6] The phosphite esters and tertiary phosphines also effect reduction:

ROOH + PR3 → OPR3 + ROH

Uses

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Precursors to epoxides

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"The single most important synthetic application of alkyl hydroperoxides is without doubt the metal-catalysed epoxidation of alkenes." In the Halcon process tert-butyl hydroperoxide (TBHP) is employed for the production of propylene oxide.[7]

Of specialized interest, chiral epoxides are prepared using hydroperoxides as reagents in the Sharpless epoxidation.[8]

The Sharpless epoxidation
The Sharpless epoxidation

Production of cyclohexanone and caprolactone

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Hydroperoxides are intermediates in the production of many organic compounds in industry. For example, the cobalt catalyzed oxidation of cyclohexane to cyclohexanone:[9]

C6H12 + O2 → (CH2)5C=O + H2O

Drying oils, as found in many paints and varnishes, function via the formation of hydroperoxides.

Hock processes

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Synthesis of cumene hydroperoxide

Compounds with allylic and benzylic C−H bonds are especially susceptible to oxygenation.[10] Such reactivity is exploited industrially on a large scale for the production of phenol by the Cumene process or Hock process for its cumene and cumene hydroperoxide intermediates.[11] Such reactions rely on radical initiators that reacts with oxygen to form an intermediate that abstracts a hydrogen atom from a weak C-H bond. The resulting radical binds O2, to give hydroperoxyl (ROO•), which then continues the cycle of H-atom abstraction.[12]

Synthesis of hydroperoxides of alkene and singlet oxygen in an Schenck ene reaction

Formation

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By autoxidation

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The most important (in a commercial sense) peroxides are produced by autoxidation, the direct reaction of O2 with a hydrocarbon. Autoxidation is a radical reaction that begins with the abstraction of an H atom from a relatively weak C-H bond. Important compounds made in this way include tert-butyl hydroperoxide, cumene hydroperoxide and ethylbenzene hydroperoxide:[7]

R−H + O2 → R−OOH


Auto-oxidation reaction is also observed with common ethers, such as diethyl ether, diisopropyl ether, tetrahydrofuran, and 1,4-dioxane. An illustrative product is diethyl ether peroxide. Such compounds can result in a serious explosion when distilled.[12] To minimize this problem, commercial samples of THF are often inhibited with butylated hydroxytoluene (BHT). Distillation of THF to dryness is avoided because the explosive peroxides concentrate in the residue.

Although ether hydroperoxide often form adventitiously (i.e. autoxidation), they can be prepared in high yield by the acid-catalyzed addition of hydrogen peroxide to vinyl ethers:[13]

C2H5OCH=CH2 + H2O2 → C2H5OCH(OOH)CH3

From hydrogen peroxide

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Many industrial peroxides are produced using hydrogen peroxide. Reactions with aldehydes and ketones yield a series of compounds depending on conditions. Specific reactions include addition of hydrogen peroxide across the C=O double bond:

R2C=O + H2O2 → R2C(OH)OOH

In some cases, these hydroperoxides convert to give cyclic diperoxides:

[R2C(O2H)]2O2 → [R2C]2(O2)2 + 2 H2O

Addition of this initial adduct to a second equivalent of the carbonyl:

R2C=O + R2C(OH)OOH → [R2C(OH)]2O2

Further replacement of alcohol groups:

[R2C(OH)]2O2 + 2 H2O2 → [R2C(O2H)]2O2 + 2 H2O

Triphenylmethanol reacts with hydrogen peroxide gives the unusually stable hydroperoxide, Ph3COOH.[14]

Naturally occurring hydroperoxides

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Many hydroperoxides are derived from fatty acids, steroids, and terpenes. The biosynthesis of these species is affected extensively by enzymes.

Illustrative biosynthetic transformation involving a hydroperoxide. Here cis-3-hexenal is generated by conversion of linolenic acid to the hydroperoxide by the action of a lipoxygenase followed by the lyase-induced formation of the hemiacetal.[15]

Inorganic hydroperoxides

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Structure of a square planar palladium hydroperoxide complex

Although hydroperoxide often refers to a class of organic compounds, many inorganic or metallo-organic compounds are hydroperoxides. One example involves sodium perborate, a commercially important bleaching agent with the formula Na2[(HO)2B]2(OO)2)]. It acts by hydrolysis to give a boron-hydroperoxide:[16]

[(HO)2B]2(OO)2)2− + 2 H2O ⇌ 2 [(HO)3B(OOH)]−

This hydrogen peroxide then releases hydrogen peroxide:

[(HO)3B(OOH)]− + H2O ⇌ B(OH)−4 + H2O2

Several metal hydroperoxide complexes have been characterized by X-ray crystallography, for example: triphenylsilicon and triphenylgermanium hydroperoxides can be obtained by reaction of initial chlorides with excess of hydrogen peroxide in presence of base.[17][18] Some form by the reaction of metal hydrides with oxygen gas:[19]

LnM−H + O2 → LnM−O−O−H (Ln refers to other ligands bound to the metal)

Some transition metal dioxygen complexes abstract H atoms (and sometimes protons) to give hydroperoxides:

LnM(O2) + H → LnMOOH

References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Hydroperoxide

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from Grokipedia
A hydroperoxide is a class of organic compounds characterized by the presence of the hydroperoxy (-OOH), consisting of an organic moiety (such as an or ) bonded to the oxygen atom of a hydroperoxy unit. These compounds can be regarded as derivatives of (H₂O₂), where one is substituted by an organic residue, resulting in the general formula ROOH. Common examples include tert-butyl hydroperoxide (t-BuOOH) and , which are notable for their role as selective oxygen-transfer agents in chemical reactions. Hydroperoxides serve as critical intermediates in processes, where they form through hydrogen abstraction from peroxyl radicals during the oxidation of hydrocarbons and other organic substrates. In , they are extensively used as oxidizing reagents, facilitating reactions such as epoxidations of alkenes, Baeyer-Villiger oxidations, and the industrial production of and acetone via the . Their reactivity stems from the weak O-O bond, which readily undergoes homolytic or heterolytic cleavage to generate reactive species like alkoxy radicals or hydroxyl anions. Despite their utility, hydroperoxides exhibit instability due to this labile linkage, often decomposing at relatively low temperatures (below 300°C) and posing risks when subjected to shock, , or contaminants such as metals. They are also key players in and combustion kinetics, contributing to the formation of secondary organic aerosols and influencing fuel oxidation pathways.

Introduction

Definition and Nomenclature

Hydroperoxides are a class of chemical compounds characterized by the general ROOH, where R represents an organic such as an alkyl or . This structure distinguishes them from peroxides, which have the ROOR with two organic s linked by an oxygen-oxygen bond. The defining is the hydroperoxy moiety (-OOH), which imparts oxidative properties to these molecules. Hydroperoxides are organic derivatives of (H₂O₂), formed by replacing one with an organic group R. In IUPAC nomenclature, hydroperoxides are named substitutively using the prefix "hydroperoxy-" attached to the parent hydride name, or as "alkane-peroxols" for acyclic structures. For example, the compound commonly known as tert-butyl hydroperoxide is systematically named 2-methylpropane-2-peroxol. Similarly, , an important industrial compound, has the IUPAC name 2-phenylpropane-2-peroxol. Common names, such as alkyl hydroperoxide, are also widely used in literature and industry for simplicity, particularly when referring to specific derivatives. The term "hydroperoxide" originates from its analogy to , reflecting the shared -OOH group and the monosubstituted derivative structure of ROOH from the parent H₂O₂ (dioxidane).

Historical Discovery

The discovery of hydroperoxides traces back to the isolation of in 1818 by French chemist Louis Jacques Thénard. Thénard prepared it by reacting with , yielding a solution he termed "eau oxygénée" (oxygenated water) due to its ability to release oxygen. This marked the first recognition of peroxides as distinct chemical entities, laying the foundation for understanding hydroperoxide chemistry. In the early , investigations into —the spontaneous oxidation of organic compounds by molecular oxygen—revealed the role of organic hydroperoxides as key intermediates. German chemist Carl Engler and his collaborators conducted pivotal studies around 1900–1910, demonstrating that of alkenes and hydrocarbons produced peroxides rather than or alone. These works established the peroxide theory of oxidation, later refined by the Bach-Engler mechanism, and confirmed the hydroperoxide structure (R-OOH) through experiments on compounds like , resolving earlier debates over cyclic formations. Engler's contributions highlighted the radical-chain nature of these processes, shifting focus from inorganic to organic variants. A major industrial milestone occurred in the 1940s with the development of the Hock process, which utilized for phenol and acetone production. German chemists Heinrich Hock and Stefan Lang reported in 1944 that air oxidation of (isopropylbenzene) yields , which undergoes acid-catalyzed rearrangement to the desired products. This process, commercialized shortly after , represented the first large-scale application of an organic hydroperoxide and underscored its synthetic utility. Hock's innovation built directly on earlier research, transforming laboratory observations into a cornerstone of modern chemical manufacturing.

Structure and Properties

Molecular Structure

Hydroperoxides contain the hydroperoxy (-OOH), where an organic residue (R) or is bonded to one oxygen atom, forming R-O-O-H. This group exhibits a characteristic linkage with an O-O length of approximately 1.45 , as observed in both and alkyl derivatives through and computational studies. The adjacent C-O bond in organic hydroperoxides is similarly around 1.45 , reflecting the single-bond character influenced by the of oxygen. The geometry of the O-O-H moiety features a bond angle of about 110°, akin to the tetrahedral arrangement in due to valence shell electron pair repulsion (, where the lone pairs on the terminal oxygen distort the ideal 109.5° angle slightly. In the gas phase, the O-O-H angle in is 94.8°. In the solid state, neutron diffraction measurements give about 102.7°. Solid-state structures of organic hydroperoxides typically show angles around 100–110°.. The electronic structure of the hydroperoxy group accounts for the relative weakness of the O-O bond, primarily due to significant repulsion between the s on the two oxygen atoms, which reduces the and lengthens the bond compared to typical single bonds. This repulsion is exacerbated by the high on oxygen, leading to bond dissociation energies around 45 kcal/mol. Structural variations arise between primary and tertiary hydroperoxides due to from the alkyl substituent. In primary hydroperoxides like ethyl hydroperoxide, the unhindered -CH₂-O-O-H chain allows for a more flexible conformation with minimal torsional strain. In contrast, tertiary examples such as tert-butyl hydroperoxide (tBuOOH) feature a bulky (CH₃)₃C- group, which introduces steric hindrance that favors a gauche conformation of the O-O-H moiety to minimize interactions, potentially slightly widening the C-O-O angle beyond 110° as predicted by calculations.

Physical Properties

Hydroperoxides are typically colorless liquids at , although those with larger or aromatic R groups, such as , may appear pale yellow. Some hydroperoxides with bulky substituents can form solids under certain conditions. Organic hydroperoxides exhibit greater volatility than the corresponding alcohols due to weaker intermolecular hydrogen bonding in the -OOH . For example, tert-butyl hydroperoxide distills at 37 °C under reduced (15 mmHg), in contrast to tert-butanol's of 82 °C at (760 mmHg). These compounds are generally soluble in organic solvents such as alcohols, ethers, and hydrocarbons, but their in decreases with increasing size of the R group. (R = H) is fully miscible with , while tert-butyl hydroperoxide (R = tert-butyl) forms stable solutions up to approximately 70 wt% in but has limited for the pure compound (around 12 wt% at 20 °C). Larger hydroperoxides like have a of 1.5 g/100 mL at 20 °C. Densities of organic hydroperoxides typically range from 0.9 to 1.1 g/cm³ at 20–25 °C and increase with the length of the alkyl chain in the R group. For instance, tert-butyl hydroperoxide has a density of 0.896 g/cm³ at 25 °C, compared to 1.029 g/cm³ for . Viscosities also rise with chain length, influenced by molecular size and hydrogen bonding; exhibits a viscosity of 1.245 mPa·s at 20 °C, higher than that of (0.890 mPa·s), while tert-butyl hydroperoxide has a viscosity of about 4 mPa·s at 20 °C.

Chemical Properties

Hydroperoxides exhibit mild acidity due to the O-H bond in the ROOH , with pKa values typically ranging from 11.5 for methyl hydroperoxide (CH₃OOH) to approximately 13 for more sterically hindered derivatives. of hydroperoxides yields the hydroperoxide anion (ROO⁻), a strong base and that plays roles in various chemical and biological processes. The O-O bond in hydroperoxides has a relatively low bond dissociation energy of 45-50 kcal/mol (190-210 kJ/mol), contributing to their thermal instability and propensity for homolytic cleavage. This weakness, compared to typical C-C or C-O bonds, underlies their reactivity in oxidation pathways. Hydroperoxides function as mild oxidants owing to the labile nature of their peroxide oxygen, enabling selective oxygen transfer in synthetic transformations such as epoxidations and alcohol oxidations under controlled conditions. These compounds are particularly sensitive to light and trace transition metals (e.g., Fe, Cu, Mn), which catalyze their decomposition via radical mechanisms, often leading to explosive hazards if not stabilized.

Formation

Autoxidation

Autoxidation is a free radical that leads to the formation of hydroperoxides (ROOH) from hydrocarbons (RH) and molecular oxygen (O₂), typically occurring under mild conditions and serving as a key pathway for oxidative degradation in organic materials. The overall simplified equation for the process is RH + O₂ → ROOH, representing the net incorporation of oxygen into the C-H bond to form the O-O-H linkage. This reaction is ubiquitous in both and , where it initiates the breakdown of alkanes, alkenes, and other substrates. The mechanism proceeds through three main stages: , , and termination. In the step, a small amount of —often from , , or added initiators—abstracts a from the , generating an alkyl radical (R•) via homolytic cleavage: RH → R• + H•. This step is rate-determining and can be accelerated by peroxides or azo compounds as initiators. follows in two coupled reactions: the alkyl radical rapidly reacts with O₂ to form a peroxyl radical (ROO•), R• + O₂ → ROO•, which then abstracts a hydrogen from another RH molecule to yield the hydroperoxide and regenerate R•, ROO• + RH → ROOH + R•. This cycle sustains the chain, with each loop producing one ROOH . Termination occurs when radicals combine, such as 2ROO• → ROOR + O₂ or ROO• + R• → ROOR, halting the chain and forming non-radical products like peroxides or alcohols. Several factors influence the rate and selectivity of . Elevated temperatures (typically 50–150°C) enhance radical formation and , while initiators like benzoyl peroxide lower the for initiation. The reaction shows high selectivity for allylic or tertiary C-H bonds due to their weaker bond dissociation energies (around 88–91 kcal/mol for allylic vs. 98 kcal/mol for primary), favoring hydroperoxide formation at these positions over secondary or primary sites. Oxygen concentration also plays a role; under high O₂ , peroxyl radical formation dominates, but low oxygen can lead to alternative pathways. Representative examples illustrate practical applications of . In the oxidation of , a liquid-phase process at 150–160°C and 10–20 bar O₂ yields cyclohexyl hydroperoxide as an intermediate, which is subsequently converted to for production, achieving selectivities up to 80–90% under optimized conditions. Similarly, the involves autoxidation of (isopropylbenzene) at 90–130°C with air, producing in yields exceeding 90%, serving as a precursor to phenol and acetone via acid-catalyzed rearrangement. These processes highlight autoxidation's efficiency in selective oxygenation, though side reactions like hydroperoxide can reduce yields if not controlled.

From Hydrogen Peroxide

One common synthetic route to organic hydroperoxides involves the of to carbonyl compounds, typically under acid or , yielding dihydroperoxides of the general formula R₂C(OOH)₂. This reaction proceeds via initial formation of a hydroperoxy hemiketal intermediate, R₂C(OH)OOH, which can further react with additional H₂O₂ to form the gem-diperoxide. Yields for this process range from 45% to 95%, depending on the catalyst, such as SnCl₄ or heteropoly acids, and reaction conditions like with 30–50% aqueous H₂O₂. A representative example is the preparation of tert-butyl hydroperoxide (TBHP), achieved by acid-catalyzed reaction of tert-butanol with 30–50% H₂O₂, using as the catalyst at moderate temperatures (40–60°C), followed by to isolate the product in 70–85% yield. The general equation for TBHP synthesis is (CH₃)₃COH + H₂O₂ → (CH₃)₃COOH + H₂O, where the alcohol acts as the precursor to the . Another approach, known as perhydrolysis, entails the acid-catalyzed addition of H₂O₂ to epoxides, resulting in ring opening to form β-hydroxy hydroperoxides. For instance, catalyzes the reaction of various epoxides with ethereal H₂O₂ at , affording β-hydroxy hydroperoxides in 70–90% yields with high favoring attack at the less substituted carbon. Similar acid-catalyzed perhydrolysis can apply to alkenes, generating hydroperoxy alcohols via , though routes are more commonly employed for controlled synthesis. These methods leverage H₂O₂ as an inexpensive, environmentally benign oxidant, enabling scalable production, although yields may be limited (sometimes below 50%) for sterically hindered substrates or without optimized catalysts.

Natural Formation

Hydroperoxides arise naturally in lipid-rich environments through the of unsaturated fatty acids, a free radical-mediated process where polyunsaturated fatty acids, such as , react with molecular oxygen to form initial alkyl radicals that propagate into peroxy radicals and ultimately yield hydroperoxides like 13-hydroperoxy-9,11-octadecadienoic acid. This occurs in biological membranes, plant oils, and animal fats, contributing to oxidative degradation akin to rancidity in stored or cellular stress in organisms. In living organisms, lipoxygenases—non-heme iron-containing dioxygenases—enzymatically catalyze the regio- and stereospecific insertion of oxygen into polyunsaturated fatty acids, producing chiral hydroperoxides such as 13S-hydroperoxy-9Z,11E-octadecadienoic acid from linoleic acid in plants and mammals. These enzymes are widespread in higher plants for jasmonate biosynthesis and in animal tissues for signaling pathways, with examples including soybean lipoxygenase-1 and human 15-lipoxygenase, which initiate hydroperoxide formation at specific allylic positions to support developmental and defensive responses. In the atmosphere, organic hydroperoxides form through the oxidation of volatile organic compounds (VOCs), primarily via reactions of organic peroxy radicals (RO₂•) with hydroperoxyl radicals (HO₂•), such as RO₂• + HO₂• → ROOH + O₂, contributing to the formation of secondary organic aerosols and serving as temporary reservoirs for atmospheric radicals. Notable natural examples include hydroperoxides in plant defense, where lipoxygenase-derived products like 9-hydroperoxy-10E,12Z-octadecadienoic acid act as precursors to volatiles in wounded tissues. In microbial contexts, certain and fungi generate organic hydroperoxides, such as those from linoleate oxidation, as byproducts of aerobic or stress responses in natural waters. Additionally, cyclic peroxides like in ambrosioides essential oil represent peroxide analogs formed via enzymatic photooxygenation, contributing to antipathogenic properties.

Reactions

Reduction

Hydroperoxides can be reduced to the corresponding alcohols using lithium aluminum hydride (LiAlH4), a strong reducing agent that cleaves the O-O bond selectively. The balanced reaction for this process is given by the equation: 4 \ceROOH+\ceLiAlH4→\ceLiAlO2+2 \ceH2O+4 \ceROH4 \ \ce{ROOH} + \ce{LiAlH4} \rightarrow \ce{LiAlO2} + 2 \ \ce{H2O} + 4 \ \ce{ROH} This reduction proceeds under mild conditions in ether solvents, yielding alcohols in high efficiency for various alkyl hydroperoxides. Other reductants include dialkyl phosphites, which react according to: \ceROOH+(RO)2P(O)H−>ROH+(RO)2P(O)OH\ce{ROOH + (RO)2P(O)H -> ROH + (RO)2P(O)OH} This method is particularly useful for analytical purposes or when avoiding metal-based reagents, as it quantitatively converts hydroperoxides to alcohols while forming phosphonic acid derivatives. (NaBH4) is employed in specific cases, such as the reduction of β-hydroxy hydroperoxides derived from photooxygenation, affording diols or triols with good yields in protic solvents like . The general mechanism for these reductions involves nucleophilic attack by the reductant on the distal oxygen of the hydroperoxide, leading to O-O bond cleavage and formation of the alcohol. The weak acidity of the O-H bond in hydroperoxides facilitates initial in some cases, enhancing reactivity. These reductions often exhibit high selectivity, preserving the at the carbon bearing the hydroperoxy group in chiral substrates, which is advantageous for stereocontrolled syntheses.

Epoxidation and Oxidation

Hydroperoxides act as effective oxygen atom donors in the metal-catalyzed epoxidation of s, where an alkyl hydroperoxide (ROOH) reacts with an to form an and the corresponding alcohol (ROH). This process typically requires transition metal catalysts such as , , or to activate the hydroperoxide, enabling stereospecific oxygen transfer while preserving the 's geometry. The reaction proceeds under mild conditions, often in organic solvents, and is widely used for synthesizing s from unfunctionalized s. A prominent variant is the Sharpless asymmetric epoxidation, which employs tert-butyl hydroperoxide (t-BuOOH) as the oxidant, along with titanium tetraisopropoxide (Ti(OiPr)4) and a chiral diethyl tartrate ligand, to produce enantiomerically enriched epoxy alcohols from allylic alcohols. This method achieves high enantioselectivity (up to >95% ) through a directed mechanism where the allylic alcohol coordinates to the titanium center, guiding the oxygen delivery from the hydroperoxide to one face of the . Developed in the early 1980s, it has become a cornerstone for asymmetric synthesis in due to its predictability and broad substrate scope. Hydroperoxides are also used in catalytic Baeyer-Villiger oxidations of ketones to esters or lactones, typically with catalysts like tin or compounds that activate the hydroperoxide for oxygen insertion. Beyond epoxidation, hydroperoxides facilitate the oxidation of sulfides to sulfoxides, often under metal-catalyzed conditions that prevent over-oxidation to sulfones. For instance, tert-butyl hydroperoxide with or catalysts selectively converts thioethers to sulfoxides in high yields, exploiting the hydroperoxide's ability to deliver a single oxygen atom. Similarly, hydroperoxides oxidize tertiary s to N-oxides, a transformation commonly achieved with t-BuOOH in the presence of group 5 or 6 metal catalysts, yielding amine oxides useful as synthetic intermediates. An industrial example is the Halcon process, where tert-butyl hydroperoxide epoxidizes to using a soluble catalyst, co-producing tert-butanol as a valuable . This method highlights the scalability of hydroperoxide-based epoxidations, operating efficiently at moderate temperatures and pressures to achieve high selectivity (>90%) for the .

Hydroperoxides undergo decomposition through various pathways, including homolytic cleavage, rearrangements, and catalyzed processes, often initiated by , metal, or /base conditions. These reactions typically break the weak O-O bond, leading to radical or ionic intermediates that propagate further transformations. Homolytic cleavage of the O-O bond in hydroperoxides (ROOH) produces alkoxy (RO•) and hydroxyl (•OH) radicals, a process favored by the relatively low bond dissociation energy of approximately 40-50 kcal/mol. This is accelerated by ions such as iron or copper, which facilitate one-electron transfer to generate the radicals. For example, in the presence of Fe(III), ROOH decomposes via a Fenton-like mechanism to initiate radical chains. A prominent rearrangement pathway occurs under , exemplified by the Hock process where alkyl hydroperoxides cleave to form a and an alcohol. In the industrial synthesis of phenol, (C6H5C(CH3)2OOH) rearranges in the presence of to yield phenol (C6H5OH) and acetone ((CH3)2C=O). This heterolytic mechanism involves of the hydroperoxide oxygen, followed by migration of an alkyl group and cleavage of the O-O bond. generally promotes decomposition to alcohols or carbonyl compounds with water, depending on the substrate structure. Base catalysis can promote decomposition via of the O-H bond, generating alkoxy oxide ions (ROO-) that may undergo further reactions depending on conditions and substrate. Uncontrolled decomposition can propagate as radical chain reactions, where initial radicals abstract hydrogens to generate new ROOH molecules, potentially leading to explosive if heat buildup is not managed.

Uses

Industrial Applications

Hydroperoxides are pivotal intermediates in several major for producing high-volume chemicals, leveraging their reactivity for efficient oxidation under mild conditions. The Hock process represents the dominant route for phenol and acetone production, accounting for the majority of global output. (isopropylbenzene) is autoxidized with air at elevated temperatures to yield , which undergoes acid-catalyzed decomposition to phenol and acetone in a 1:1 molar ratio. This method supplies over 95% of the world's phenol, with global production capacity reaching 16.06 million metric tons per annum in 2023. In the manufacture of , a key precursor for nylon-6 and , is subjected to air oxidation to form cyclohexyl hydroperoxide as the primary intermediate. This hydroperoxide is then thermally rearranged or catalytically converted to a mixture of and (known as KA oil), from which is separated and further processed. The oxidation route dominates, comprising about 65% of global capacity, estimated at 4.97 million metric tons in 2024. The Halcon process employs tert-butyl hydroperoxide (TBHP) for the selective epoxidation of to , an essential building block for polyether polyols and propylene glycols. is oxidized to TBHP using air, and the purified hydroperoxide reacts with over a soluble catalyst to produce and tert-butanol as a , which can be dehydrated to isobutene for recycling. This technology supports a portion of the global propylene oxide market, which totals approximately 10.1 million metric tons in 2024. These processes derive economic advantages from molecular oxygen as an inexpensive, abundant feedstock derived from air, minimizing raw material costs and enabling large-scale operations with high throughput. Nonetheless, purification of the hydroperoxide streams poses significant challenges due to low selectivity in autoxidation steps, resulting in mixtures contaminated with alcohols, ketones, and dicarboxylic acids that necessitate energy-intensive distillation and extraction to achieve commercial purity and mitigate decomposition risks.

Synthetic Applications

Hydroperoxides, particularly tert-butyl hydroperoxide (t-BuOOH), play a pivotal role in enantioselective synthesis, most notably through the Sharpless asymmetric epoxidation developed in 1980. This method employs t-BuOOH as the terminal oxidant in conjunction with titanium tetraisopropoxide and a chiral diethyl tartrate ligand to achieve high enantioselectivity in the epoxidation of allylic alcohols, enabling the predictable formation of epoxy alcohols with ee values often exceeding 90%. The reaction's reliability and broad substrate scope have made it indispensable for constructing complex chiral building blocks in natural product synthesis, contributing to K. Barry Sharpless's share of the 2001 Nobel Prize in Chemistry for work on chirally catalyzed oxidation reactions. In , hydroperoxides facilitate precise oxidations that enhance stereocontrol and efficiency. For instance, t-BuOOH has been utilized in the Jacobsen group's convergent assembly of chiral components for the of FR901464, a spliceostatin antitumor agent precursor, where it serves as an oxidant in the preparation of key fragments with high . This application exemplifies how hydroperoxides enable hydrolytic kinetic resolution-like processes in multistep sequences, allowing access to enantiopure intermediates essential for pharmaceutical targets. Such uses underscore hydroperoxides' versatility in enabling regioselective and stereospecific transformations within intricate synthetic routes. Hydroperoxides offer advantages in green chemistry by serving as safer, less reactive alternatives to traditional peracids like m-chloroperoxybenzoic acid (mCPBA), which can pose handling risks due to their instability and potential for explosive decomposition. Unlike peracids, alkyl hydroperoxides such as t-BuOOH exhibit greater thermal stability and solubility in organic solvents, reducing byproduct formation and facilitating milder conditions that align with sustainable synthesis principles, including and waste minimization. Recent advancements (2023–2025) have integrated hydroperoxides into continuous-flow chemistry platforms, enhancing safety by minimizing accumulation of reactive intermediates and enabling precise control over exothermic reactions. For example, systems have been developed for the on-demand generation and use of t-BuOOH in epoxidations and other oxidations, achieving high yields while mitigating hazards associated with . These flow-based approaches support scalable synthesis for pharmaceutical applications, with demonstrated long-term stability in producing epoxides from alkenes using hydroperoxide oxidants.

Inorganic Hydroperoxides

Examples and Structure

Hydrogen peroxide (H₂O₂) is the simplest inorganic hydroperoxide, characterized by a non-linear, skewed structure with the formula H-O-O-H, where the O-O bond length is approximately 1.47 Å, longer than a typical O-O single bond due to repulsion between the lone pairs on the oxygen atoms. Sodium perborate (NaBO₃·4H₂O) serves as another key example, functioning as an adduct that hydrolyzes in water to release hydrogen peroxide through boron-oxygen-oxygen bonds, with the tetrahydrate form exhibiting a crystalline structure where peroxy groups are integrated into the borate framework. In solution, it participates in an equilibrium involving hydroperoxyborate species, such as [(HO)2B(O2)2]2−+2H2O⇌2[(HO)3B(OOH)]−[(HO)_2B(O_2)_2]^{2-} + 2H_2O \rightleftharpoons 2[(HO)_3B(OOH)]^{-}, which underscores its role as a stable source of active oxygen. Metal hydroperoxide complexes represent a significant class in biological systems, exemplified by the [Fe(III)(OOH)] moiety in enzymes like , where the hydroperoxide ligand coordinates to the ferric iron center via the terminal oxygen, facilitating oxygen transfer reactions. Other examples include hydroperoxides like sodium hydroperoxide (NaOOH) and potassium hydroperoxide (KOOH), which are highly reactive and unstable solids prepared at low temperatures. While simple inorganic hydroperoxides like H₂O₂ exhibit good stability, those involving metal centers or ionic salts often show reduced stability due to their polar and ionic nature, promoting decomposition under shock or heat, though generally less explosive than many organic hydroperoxides.

Preparation

The primary inorganic hydroperoxide, (H₂O₂), is predominantly produced on an industrial scale via the , a cyclic oxidation-reduction method involving the of with hydrogen gas over a catalyst to form anthrahydroquinone, followed by oxidation with atmospheric oxygen to regenerate the and liberate H₂O₂. This process, which accounts for over 95% of global H₂O₂ production, operates in a mixture and yields concentrations up to 40% H₂O₂ after extraction and purification, emphasizing its efficiency in avoiding direct H₂/O₂ mixing to mitigate explosion risks. An alternative electrolytic method involves the of dilute or solutions using electrodes, where anodic oxidation produces peroxodisulfate ions (S₂O₈²⁻) that hydrolyze to H₂O₂ upon , though this approach is less common today due to higher energy demands compared to the anthraquinone route. Sodium perborate (NaBO₃·nH₂O), a key solid inorganic hydroperoxide derivative, is synthesized by reacting sodium metaborate (NaBO₂) with hydrogen peroxide in an aqueous solution at controlled pH and temperature, typically forming the tetrahydrate (n=4) that can be dehydrated to the monohydrate for stability. This straightforward neutralization process leverages H₂O₂ as the active oxygen source, with the borate matrix enhancing solubility and storage properties, and yields products containing about 10% available oxygen. Metal hydroperoxides, such as those of s (e.g., Ti(OOH)₂ or alkylperoxo complexes), are prepared through the insertion of molecular oxygen (O₂) into metal-hydride bonds or by reaction with existing peroxides, often under mild conditions to form M-OOH where M is a late like or . For instance, O₂ reacts with metal hydrides (M-H) to yield hydroperoxide intermediates via , which can be isolated or further transformed, providing a route to catalytically active in oxidation chemistry. These methods are typically conducted in inert atmospheres to prevent side reactions, with structural confirmation via revealing the η²-OOH coordination. Recent advancements (2023–2025) in sustainable H₂O₂ production focus on electrocatalytic two-electron (2e⁻ ORR) using carbon-based or metal-doped catalysts in alkaline or neutral electrolytes, achieving selectivities over 90% and production rates up to approximately 2.5 mol g⁻¹ h⁻¹ cat under ambient conditions, as demonstrated with boron-doped or nitrogen-rich carbon materials that favor the O₂ to H₂O₂ pathway over full reduction to . These approaches, integrated with renewable , offer a greener alternative to traditional processes by utilizing air-sourced O₂ and avoiding organic solvents or high-pressure . Further progress as of 2025 includes catalysts like Ni-BTA with enhanced performance in neutral media.

Safety and Handling

Hazards

Hydroperoxides, particularly organic variants such as tert-butyl hydroperoxide (TBHP), pose significant and hazards due to their instability and strong oxidizing properties. Pure or highly concentrated forms (>90%) of TBHP are shock-sensitive and can decompose violently upon heating, impact, or , potentially leading to or rapid pressure buildup in confined spaces. Even solutions between 72% and 90% concentration are classified as highly reactive and require stabilization and specific packaging for shipment, as they form explosive vapor-air mixtures above 43°C; concentrations above 90% are forbidden for . This instability stems from their tendency to undergo exothermic , releasing oxygen that intensifies fires involving nearby combustibles. Corrosivity is a major concern with hydroperoxides, especially in concentrations exceeding 10%, where they can cause severe burns to the skin and eyes upon contact. TBHP solutions above this threshold lead to rapid tissue damage, including blistering, ulceration, and potential permanent eye injury or blindness. Inhalation of vapors or mists irritates the respiratory tract, causing coughing, shortness of breath, and corrosive effects on mucous membranes, with higher concentrations exacerbating risks of pulmonary edema. Toxicity profiles of hydroperoxides include acute and potential chronic effects; for instance, TBHP exhibits mutagenic potential . is particularly dangerous, with an oral LD50 of 50–500 mg/kg for TBHP indicating severe that can be fatal in significant doses, leading to gastrointestinal , systemic absorption, and organ failure. Reactivity hazards arise from hydroperoxides' incompatibility with reducing agents, metals (e.g., iron, ), and organic materials, which can trigger runaway exothermic reactions or spontaneous . Contact with such substances may accelerate oxygen release, escalating to fires or explosions, as seen in reactions with contaminants that catalyze breakdown. This mirrors their general instability, where trace impurities promote violent exothermic events.

Storage and Precautions

Hydroperoxides, both inorganic and organic, require specific storage conditions to minimize and maintain stability. Inorganic hydroperoxides such as (H₂O₂) should be stored in cool, dry, well-ventilated areas below 40°C, using original vented containers made of passivated , , or to allow for gas release and prevent buildup. Organic hydroperoxides, exemplified by tert-butyl hydroperoxide, must be kept in tightly closed containers in locked, cool, dark locations, often refrigerated or below their self-accelerating temperature (SADT) as specified in safety data sheets, with quantities limited per NFPA 400 guidelines to reduce risk. Stabilizers are commonly added; for instance, at approximately 200 ppm is used in some H₂O₂ solutions to inhibit , while organic hydroperoxides may incorporate phlegmatizers or diluents for desensitization. Safe handling practices emphasize (PPE) and compatibility controls. Personnel should wear chemical-resistant gloves (e.g., , , or PVC), safety goggles or face shields, and protective clothing such as flame-retardant lab coats or suits to prevent and . Avoid contact with metals (e.g., , iron), alkalies, reducing agents, or contaminants, as these can accelerate decomposition; use compatible materials like 316 or PTFE for equipment. For H₂O₂, dilution with is recommended during transfer to reduce concentration and volatility, and unused material should never be returned to storage containers. Regulatory limits include OSHA's (PEL) of 1 ppm (1.4 mg/m³) for H₂O₂ over an 8-hour workday, with and monitoring required in handling areas. Emergency response protocols focus on , neutralization, and medical aid. Facilities must provide stations, showers, and adequate supplies for immediate flushing of exposures. For spills, evacuate the area, absorb with inert materials like or sand, and neutralize residues using solution to decompose the hydroperoxide; dilute large spills with copious while avoiding confinement. Collected waste should be disposed of as hazardous material per local regulations, and professional emergency services contacted for incidents involving fire or large quantities.

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