Trioxidane

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

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.

Preparation

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). 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. In acetone-''d''₆ at −20 Â°C, the characteristic ¹H NMR signal of trioxidane could be observed at a chemical shift of 13.1 ppm. Solutions of hydrogen trioxide in diethyl ether can be safely stored at −20 Â°C for as long as a week.

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.

Structure

In 1970-75, Giguère et al. observed infrared and Raman spectra of dilute aqueous solutions of trioxidane. 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. Various dimeric and trimeric forms also seem to exist.

There is a trend of increasing gas-phase acidity and corresponding p''K''_a as the number of oxygen atoms in the chain increases in HO_''n''H structures (''n''=1,2,3).

Reactions

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. 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.

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 type of molecules can start in interstellar space. A 2022 publication suggested the possibility of the presence of detectable concentrations of polyoxides in the atmosphere.

See also

* Molozonide

References

{{Hydrides by group

Inorganic compounds
Oxides
Polyoxides
Hydrogen compounds