Selenic acid
| |||
| Names | |||
|---|---|---|---|
| IUPAC name
Selenic(VI) acid
| |||
| Other names
Selenic acid
| |||
| Identifiers | |||
3D model (JSmol)
|
|||
| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.029.072 | ||
| KEGG | |||
PubChem CID
|
|||
| RTECS number |
| ||
CompTox Dashboard (EPA)
|
|||
| |||
| |||
| Properties | |||
| H 2SeO 4 | |||
| Molar mass | 144.9734 g/mol | ||
| Appearance | Colorless deliquescent crystals | ||
| Density | 2.95 g/cm3, solid | ||
| Melting point | 58 °C (331 K) | ||
| Boiling point | 260 °C (533 K) (decomposes) | ||
| 130 g/100 mL (30 °C) | |||
| Acidity (pKa) | similar to H 2SO 4 | ||
Refractive index (nD)
|
1.5174 (D-line, 20 °C) | ||
| Structure | |||
| tetrahedral at Se | |||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards
|
Corrosive, highly toxic | ||
| NFPA 704 (fire diamond) | |||
| Related compounds | |||
Other anions
|
selenious acid hydroselenic acid hydrogen selenide | ||
Other cations
|
sodium selenate | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |||
Selenic acid is the chemical compound with the formula H
2SeO
4. It is an oxoacid of selenium, and its structure is more accurately described as (HO)
2SeO
2.
As predicted by VSEPR theory, the selenium center is tetrahedral, with a Se–O bond length of 161 pm.[1] In the solid state, it crystallizes in an orthorhombic structure.[2]
Preparation
Due to the instability of selenium trioxide, it is impractical to synthesize selenic acid by dissolving selenium trioxide in water, unlike sulfuric acid, which can be synthesized by dissolving sulfur trioxide in water.[1] Instead, it must be prepared by oxidising selenium compounds in lower oxidation states.
One method of preparing selenic acid is by the oxidation of selenium dioxide with hydrogen peroxide:
- SeO
2 + H
2O
2 → H
2SeO
4
To obtain the anhydrous acid as a crystalline solid, the resulting solution is evaporated at temperatures below 140 °C (413 K; 284 °F) in a vacuum.[3]
Selenic acid may also be prepared by the oxidation of selenous acid (H
2SeO
3) with halogens, such as chlorine or bromine, or with potassium permanganate.[4] However, using chlorine or bromine as the oxidising agent also produces hydrochloric or hydrobromic acid as a side-product, which needs to be removed from the solution since they can reduce the selenic acid to selenous acid.[5]
Another method of preparing selenic acid is by the oxidation of elemental selenium in water suspension by chlorine:[4]
- Se + 4 H
2O + 3 Cl
2 → H
2SeO
4 + 6 HCl
Chemistry
Like sulfuric acid, selenic acid is a strong acid that is hygroscopic and extremely soluble in water. Concentrated solutions are viscous. Crystalline mono- and di-hydrates are known.[4] The monohydrate melts at 26°C, and the dihydrate melts at −51.7°C.[1]
Selenic acid is a stronger oxidiser than sulfuric acid,[4] capable of liberating chlorine from chloride ions, being reduced to selenous acid in the process:
- H
2SeO
4 + 2 H+
+ 2 Cl−
→ H
2SeO
3 + H
2O + Cl
2
It decomposes above 200°C, liberating oxygen gas and being reduced to selenous acid:[4]
- 2 H
2SeO
4 → 2 H
2SeO
3 + O
2
Selenic acid reacts with barium salts to precipitate BaSeO
4, analogous to the sulfate. In general, selenate salts resemble sulfate salts, but are more soluble. Many selenate salts have the same crystal structure as the corresponding sulfate salts.[1]
Treatment of selenic acid with fluorosulfuric acid gives the dioxydifluoride (b.p. −8.4 °C):[3]
- H
2SeO
4 + 2 HO
3SF → SeO
2F
2 + 2 H
2SO
4
Hot, concentrated selenic acid is capable of dissolving gold, forming a reddish-yellow solution of gold(III) selenate:[6]
- 2 Au + 6 H
2SeO
4 → Au
2(SeO
4)
3 + 3 H
2SeO
3 + 3 H
2O
Applications
Used as a reagent for alkaloids and as an oxidizing agent.
References
- ^ a b c d Don M. Yost (2007). Systematic Inorganic Chemistry. READ BOOKS. pp. 343–346. ISBN 1-4067-7302-6.
- ^ Mathias S. Wickleder (2007). Francesco A. Devillanova (ed.). Handbook of chalcogen chemistry: new perspectives in sulfur, selenium and tellurium. Royal Society of Chemistry. p. 353. ISBN 0-85404-366-7.
- ^ a b Seppelt, K. “Selenoyl difluoride” Inorganic Syntheses, 1980, volume XX, pp. 36-38. ISBN 0-471-07715-1. The report describes the synthesis of selenic acid.
- ^ a b c d e Anil Kumar De (2003). A Text Book of Inorganic Chemistry. New Age International. pp. 543–545. ISBN 81-224-1384-6.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1021/ja01683a005, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1021/ja01683a005instead. - ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1021/ja02018a005, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1021/ja02018a005instead.