 | |
 | |
| Names | |
|---|---|
| Preferred IUPAC name Trioxidane (only preselected name)[1] | |
| Systematic IUPAC name Dihydrogen trioxide | |
| Other names Hydrogen trioxide | |
| Identifiers | |
3D model (JSmol) | |
| ChEBI | |
| ChemSpider |
|
| 200290 | |
| |
| |
| Properties | |
| H2O3 | |
| Molar mass | 50.013 g·mol−1 |
| Related compounds | |
Related compounds | |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Trioxidane (systematically nameddihydrogen trioxide,[2][3]), also calledhydrogen trioxide[4][5] is aninorganic compound with thechemical formulaH2O3 orO(OH)2 (can be also written asH[O]3H or[H(μ-O3)H]). It is one of the unstablehydrogen polyoxides.[4] In aqueous solutions, trioxidane decomposes to form water andsinglet 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.[2]
Trioxidane can be obtained in small, but detectable, amounts in reactions ofozone andhydrogen peroxide, or by theelectrolysis of water. Larger quantities have been prepared by the reaction of ozone with organicreducing agents at low temperatures in a variety of organic solvents, such as theanthraquinone process. It is also formed during the decomposition of organic hydrotrioxides (ROOOH).[3] Alternatively, trioxidane can be prepared by reduction of ozone with1,2-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 themethyltrioxorhenium(VII) catalyst.[5] In acetone-d6 at −20 °C, the characteristic1H NMR signal of trioxidane could be observed at achemical 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]
In 1970–75,Giguère et al. observedinfrared andRaman spectra of dilute aqueous solutions of trioxidane.[4] In 2005, trioxidane was observed experimentally bymicrowave spectroscopy in a supersonic jet. The molecule exists in a skewed structure, with an oxygen–oxygen–oxygen–hydrogendihedral angle of 81.8°. The oxygen–oxygenbond lengths of 142.8picometer are slightly shorter than the 146.4 pm oxygen–oxygen bonds inhydrogen peroxide.[7] Various dimeric and trimeric forms also seem to exist.
There is a trend of increasinggas-phase acidity and corresponding pKa as the number of oxygen atoms in the chain increases in HOnH structures (n=1,2,3).[8]
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 formsulfoxides, but little else is known of its reactivity.
Recent research found that trioxidane is the active ingredient responsible for theantimicrobial properties of the well knownozone/hydrogen peroxide mix. Because these two compounds are present in biological systems as well it is argued that anantibody in the human body can generate trioxidane as a powerfuloxidant 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 ininterstellar space.[7] A 2022 publication suggested the possibility of the presence of detectable concentrations of polyoxides in the atmosphere.[11]
