Chemical compound

Selenium dioxide is the chemical compound with the formula SeO2. This colorless solid is one of the most frequently encountered compounds of selenium.

Properties

Solid SeO2 is a one-dimensional polymer, the chain consisting of alternating selenium and oxygen atoms. Each Se atom is pyramidal and bears a terminal oxide group. The bridging Se-O bond lengths are 179 pm and the terminal Se-O distance is 162 pm.[3] The relative stereochemistry at Se alternates along the polymer chain (syndiotactic). In the gas phase selenium dioxide is present as dimers and other oligomeric species, at higher temperatures it is monomeric.[4] The monomeric form adopts a bent structure very similar to that of sulfur dioxide with a bond length of 161 pm.[4] The dimeric form has been isolated in a low temperature argon matrix and vibrational spectra indicate that it has a centrosymmetric chair form.[3] Dissolution of SeO2 in selenium oxydichloride give the trimer [Se(O)O]3.[4] Monomeric SeO2 is a polar molecule, with the dipole moment of 2.62 D [5] pointed from the midpoint of the two oxygen atoms to the selenium atom.

The solid sublimes readily. At very low concentrations the vapor has a revolting odor, resembling decayed horseradishes. At higher concentrations the vapor has an odor resembling horseradish sauce and can burn the nose and throat on inhalation. Whereas SO2 tends to be molecular and SeO2 is a one-dimensional chain, TeO2 is a cross-linked polymer.[3]

SeO2 is considered an acidic oxide: it dissolves in water to form selenous acid.[4] Often the terms selenous acid and selenium dioxide are used interchangeably. It reacts with base to form selenite salts containing the SeO2−
3
anion. For example, reaction with sodium hydroxide produces sodium selenite:

SeO2 + 2 NaOH → Na2SeO3 + H2O

Preparation

Selenium dioxide is prepared by oxidation of selenium by burning in air or by reaction with nitric acid or hydrogen peroxide, but perhaps the most convenient preparation is by the dehydration of selenous acid.

2 H2O2 + Se → SeO2 + 2 H2O
3 Se + 4 HNO3 + H2O → 3 H2SeO3 + 4 NO
H2SeO3 ⇌ SeO2 + H2O

Occurrence

The natural form of selenium dioxide, downeyite, is a very rare mineral. It is found in only a very few burning coal dumps.[6]

Uses

Organic synthesis

SeO2 is an important reagent in organic synthesis. Oxidation of paraldehyde (acetaldehyde trimer) with SeO2 gives glyoxal[7] and the oxidation of cyclohexanone gives cyclohexane-1,2-dione.[8] The selenium starting material is reduced to selenium, and precipitates as a red amorphous solid which can easily be filtered off.[8] This type of reaction is called a Riley oxidation. It is also renowned as a reagent for "allylic" oxidation,[9] a reaction that entails the following conversion

Allylic oxidation

This can be described more generally as;

R2C=CR'-CHR"2 + [O] → R2C=CR'-C(OH)R"2

where R, R', R" may be alkyl or aryl substituents.

Selenium dioxide can also be used to synthesize 1,2,3-selenadiazoles from acylated hydrazone derivatives.[10]

As a colorant

Selenium dioxide imparts a red color to glass. It is used in small quantities to counteract the color due to iron impurities and so to create (apparently) colorless glass. In larger quantities, it gives a deep ruby red color.

Selenium dioxide is the active ingredient in some cold-bluing solutions.

It was also used as a toner in photographic developing.

Safety

Selenium is an essential element, but ingestion of more than 5 mg/day leads to nonspecific symptoms.[11]

References

  1. ^ http://www.integrachem.com/msds/S138_26294_101.pdf[bare URL PDF]
  2. ^ "Selenium compounds (as Se)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  3. ^ a b c Handbook of Chalcogen Chemistry: New Perspectives in Sulfur, Selenium and Tellurium, Franceso A. Devillanova, Royal Society of Chemistry, 2007, ISBN 9780854043668
  4. ^ a b c d Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
  5. ^ Takeo, Harutoshi; Hirota, Eizi; Morino, Yonezo (1972). "Third-order potential constants and dipole moment of SeO2 by microwave spectroscopy". Journal of Molecular Spectroscopy. 41 (2): 420–422. Bibcode:1972JMoSp..41..420T. doi:10.1016/0022-2852(72)90216-0. ISSN 0022-2852.
  6. ^ Finkelman, Robert B.; Mrose, Mary E. (1977). "Downeyite, the first verified natural occurrence of SeO2" (PDF). American Mineralogist. 62: 316–320.
  7. ^ Ronzio, A. R.; Waugh, T. D. (1955). "Glyoxal Bisulfite". Organic Syntheses.{{cite journal}}: CS1 maint: multiple names: authors list (link); Collective Volume, vol. 3, p. 438
  8. ^ a b Hach, C. C. Banks, C. V.; Diehl, H. (1963). "1,2-Cyclohexanedione Dioxime". Organic Syntheses.{{cite journal}}: CS1 maint: multiple names: authors list (link); Collective Volume, vol. 4, p. 229
  9. ^ Coxon, J. M.; Dansted, E.; Hartshorn, M. P. (1988). "Allylic Oxidation with Hydrogen Peroxide–Selenium Dioxide: trans-Pinocarveol". Organic Syntheses.{{cite journal}}: CS1 maint: multiple names: authors list (link); Collective Volume, vol. 6, p. 946
  10. ^ Lalezari, Iradj; Shafiee, Abbas; Yalpani, Mohamed (1969). "A novel synthesis of selenium heterocycles: substituted 1,2,3-selenadiazoles". Tetrahedron Letters. 10 (58): 5105–5106. doi:10.1016/S0040-4039(01)88895-X.
  11. ^ Bernd E. Langner "Selenium and Selenium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a23_525

External links