Nitrous acid

Nitrous acid (molecular formula HNO
2) is a weak and monoprotic acid known only in solution, in the gas phase and in the form of nitrite (NO−
2) salts.[1] Nitrous acid is used to make diazonium salts from amines. The resulting diazonium salts are reagents in azo coupling reactions to give azo dyes.

Not to be confused with nitric acid.
Nitrous acid
Names
Preferred IUPAC name
Nitrous acid
Systematic IUPAC name
Hydroxidooxidonitrogen
Identifiers
3D model (JSmol)
3DMet
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.029.057
EC Number
  • 231-963-7
983
KEGG
MeSH Nitrous+acid
UNII
Properties
HNO2
Molar mass 47.013 g/mol
Appearance Pale blue solution
Density Approx. 1 g/ml
Melting point Only known in solution or as gas
Acidity (pKa) 3.15
Conjugate base Nitrite
Hazards
NFPA 704 (fire diamond)
Flash point Non-flammable
Related compounds
Other anions
 Nitric acid
Other cations
 Sodium nitrite
Potassium nitrite
Ammonium nitrite
Related compounds
 Dinitrogen trioxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

Structure

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

Preparation

Nitrous acid is usually generated by acidification of aqueous solutions of sodium nitrite with a mineral acid. The acidification is usually conducted at ice temperatures, and the HNO2 is consumed in situ.[2][3] Free nitrous acid is unstable and decomposes rapidly.

Nitrous acid can also be produced by dissolving dinitrogen trioxide in water according to the equation

N2O3 + H2O → 2 HNO2

Reactions

Nitrous acid is the main chemphore in the Liebermann reagent, used to spot-test for alkaloids.

Decomposition
See also: Dinitrogen trioxide

Gaseous nitrous acid, which is rarely encountered, 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:[4]

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

The nitric oxide can subsequently be re-oxidized by air to nitric acid, making the overall reaction:

2 HNO2 + O2 → 2 HNO3

Reduction

With I− and Fe2+ ions, NO is formed:[5]

2 KNO2 + 2 KI + 2 H2SO4 → I2 + 2 NO + 2 H2O + 2 K2SO4
2 KNO2 + 2 FeSO4 + 2 H2SO4 → Fe2(SO4)3 + 2 NO + 2 H2O + K2SO4

With Sn2+ ions, N2O is formed:

2 KNO2 + 6 HCl + 2 SnCl2 → 2 SnCl4 + N2O + 3 H2O + 2 KCl

With SO2 gas, NH2OH is formed:

2 KNO2 + 6 H2O + 4 SO2 → 3 H2SO4 + K2SO4 + 2 NH2OH

With Zn in alkali solution, NH3 is formed:

5 H2O + KNO2 + 3 Zn → NH3 + KOH + 3 Zn(OH)2

With N
2H+
5, HN3, and subsequently, N2 gas is formed:

HNO2 + [N2H5]+ → HN3 + H2O + H3O+
HNO2 + HN3 → N2O + N2 + H2O

Oxidation by nitrous acid has a kinetic control over thermodynamic control, this is best illustrated that dilute nitrous acid is able to oxidize I− to I2, but dilute nitric acid cannot.

I2 + 2 e− ⇌ 2 I−   Eo = +0.54 V
NO−
3 + 3 H+ + 2 e− ⇌ HNO2 + H2O   Eo = +0.93 V
HNO2 + H+ + e− ⇌ NO + H2O   Eo = +0.98 V

It can be seen that the values of Eo
cell for these reactions are similar, but nitric acid is a more powerful oxidizing agent. Base on the fact that dilute nitrous acid can oxidize iodide into iodine, it can be deduced that nitrous is a faster, rather than a more powerful, oxidizing agent than dilute nitric acid.[5]

Organic chemistry

Nitrous acid is used to prepare diazonium salts:

HNO2 + ArNH2 + H+ → ArN+
2 + 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.[6] 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:[7] It is mainly blue in colour

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)2CHCH2CH2OH + HNO2 → (CH3)2CHCH2CH2ONO + H2O

The carcinogens called nitrosamines are produced, usually not intentionally, by the reaction of nitrous acid with secondary amines:

HNO2 + R2NH → R2N-NO + H2O

Atmosphere of the Earth

Nitrous acid is involved in the ozone budget of the lower atmosphere, the troposphere. The heterogeneous reaction of nitric oxide (NO) and water produces nitrous acid. When this reaction takes place on the surface of atmospheric aerosols, the product readily photolyses to hydroxyl radicals.[8][9]

See also

References
  1. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 462.
  2. Y. Petit, M. Larchevêque (1998). "Ethyl Glycidate from (S)-Serine: Ethyl (R)-(+)-2,3-Epoxypropanoate". Org. Synth. 75: 37. doi:10.15227/orgsyn.075.0037.CS1 maint: uses authors parameter (link)
  3. Adam P. Smith, Scott A. Savage, J. Christopher Love, Cassandra L. Fraser (2002). "Synthesis of 4-, 5-, and 6-methyl-2,2'-bipyridine by a Negishi Cross-coupling Strategy: 5-methyl-2,2'-bipyridine". Org. Synth. 78: 51. doi:10.15227/orgsyn.078.0051.CS1 maint: uses authors parameter (link)
  4. 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.
  5. Catherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 15: The group 15 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 449. ISBN 978-0-13-175553-6.
  6. Clarke, H. T.; Kirner, W. R. "Methyl Red" Organic Syntheses, Collected Volume 1, p.374 (1941). "Archived copy" (PDF). Archived from the original (PDF) on 2007-09-30. Retrieved 2007-07-26.CS1 maint: archived copy as title (link)
  7. Prudent practices in the laboratory: handling and disposal of chemicals. Washington, D.C.: National Academy Press. 1995. doi:10.17226/4911. ISBN 978-0-309-05229-0.
  8. Spataro, F; Ianniello, A (November 2014). "Sources of atmospheric nitrous acid: state of the science, current research needs, and future prospects". Journal of the Air & Waste Management Association. 64 (11): 1232–1250. doi:10.1080/10962247.2014.952846. PMID 25509545.
  9. Anglada, Josef M.; Solé, Albert (November 2017). "The Atmospheric Oxidation of HONO by OH, Cl, and ClO Radicals". The Journal of Physical Chemistry A. 121 (51): 9698–9707. Bibcode:2017JPCA..121.9698A. doi:10.1021/acs.jpca.7b10715. PMID 29182863.
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