Nitrous acid

Nitrous acid
Preferred IUPAC name Nitrous acid
Systematic name Hydroxidooxidonitrogen
Identifiers
CAS number	7782-77-6 Y
PubChem	24529 Y
ChemSpider	22936 Y
EC number	231-963-7
KEGG	C00088 Y
MeSH	Nitrous+acid
ChEBI	CHEBI:25567 Y
ChEMBL	CHEMBL1161681 N
Gmelin Reference	983
3DMet	B00022
Jmol-3D images	Image 1
SMILES O=NO
InChI InChI=1S/HNO2/c2-1-3/h(H,2,3) Y Key: IOVCWXUNBOPUCH-UHFFFAOYSA-N Y
Properties
Molecular formula	HNO2
Molar mass	47.013 g/mol
Appearance	Pale blue solution
Density	Approx. 1 g/ml
Melting point	Only known in solution
Acidity (pKa)	3.398
Hazards
EU Index	Not listed
Flash point	Non-flammable
Related compounds
Other anions	Nitric acid
Other cations	Sodium nitrite Potassium nitrite Ammonium nitrite
Related compounds	Dinitrogen trioxide
N (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Nitrous acid (molecular formula HNO₂) is a weak and monobasic acid known only in solution and in the form of nitrite salts.

Nitrous acid is used to make diazides from amines; this occurs by nucleophilic attack of the amine onto the nitrite, reprotonation by the surrounding solvent, and double-elimination of water. The diazide can then be liberated to give a carbene or carbenoid.

Structure

In the gas phase, the planar nitrous acid molecule can adopt both a cis and a trans form. The trans form predominates at room temperature, and IR measurements indicate it is more stable by around 2.3 kJ mol⁻¹.^[1]

dimensions of the trans form (from the microwave spectrum)	model of the trans form	cis form

Preparation

When cold, dilute solutions of nitrite ion, NO₂⁻ are carefully acidified, a light blue solution of nitrous acid is produced. Free nitrous acid is unstable and decomposes rapidly.

Decomposition

See also: Dinitrogen trioxide

In anything other than very dilute, cold solutions, nitrous acid rapidly decomposes into nitrogen dioxide, nitric oxide, and water:

2 HNO₂ → NO₂ + NO + H₂O

Nitrogen dioxide disproportionates into nitric acid and nitrous acid in aqueous solution:^[2]

2 NO₂ + H₂O → HNO₃ + HNO₂

In warm or concentrated solutions, the overall reaction amounts to production of nitric acid, water, and nitric oxide:

3 HNO₂ → HNO₃ + 2 NO + H₂O

Chemistry

Nitrous acid is used to prepare diazonium salts:

HNO₂ + ArNH₂ + H⁺ → ArN₂⁺ + 2 H₂O

where Ar is an aryl group.

Such salts are widely used in organic synthesis, e.g., for the Sandmeyer reaction and in the preparation azo dyes, brightly-colored compounds that are the basis of a qualitative test for anilines.^[3] Nitrous acid is used to destroy toxic and potentially-explosive sodium azide. For most purposes, nitrous acid is usually formed in situ by the action of mineral acid on sodium nitrite:^[4]

NaNO₂ + HCl → HNO₂ + NaCl

2 NaN₃ + 2 HNO₂ → 3 N₂ + 2 NO + 2 NaOH

Reaction with two α-hydrogen atoms in ketones creates oximes, which may be further oxidized to a carboxylic acid, or reduced to form amines. This process is used in the commercial production of adipic acid.

Nitrous acid reacts rapidly with aliphatic alcohols to produce alkyl nitrites, which are potent vasodilators:

(CH₃)₂CH-CH₂-CH₂-OH + HNO₂ → (CH₃)₂CH-CH₂-CH₂-ONO + H₂O

Atmosphere of the earth

Nitrous acid is involved in the ozone budget of the lower atmosphere: the troposphere. The heterogeneous reaction of nitrogen monoxide (NO) and water produces nitrous acid. When this reaction takes place on the surface of atmospheric aerosols, product readily photolyses to hydroxyl radicals.

See also

• Demjanov rearrangement
• Nitric acid (HNO₃)
• Tiffeneau-Demjanov rearrangement

References

1. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0080379419.  p. 462
2. ^ Kameoka, Yohji; Pigford, Robert (February 1977). "Absorption of Nitrogen Dioxide into Water, Sulfuric Acid, Sodium Hydroxide, and Alkaline Sodium Sulfite Aqueous". Ind. Eng. Chem. Fundamen. 16 (1): 163–169. doi:10.1021/i160061a031. 
3. ^ Clarke, H. T.; Kirner, W. R. "Methyl Red" Organic Syntheses, Collected Volume 1, p.374 (1941). http://www.orgsyn.org/orgsyn/pdfs/CV1P0374.pdf
4. ^ Prudent practices in the laboratory: handling and disposal of chemicals. Washington, D.C.: National Academy Press. 1995. ISBN 0309052297. http://books.nap.edu/openbook.php?record_id=4911&page=165.