Nitrous acid

Not to be confused with nitric acid.
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Nitrous acid
Nitrous acid
Names
IUPAC name
Nitrous acid[1]
Identifiers
3D model (JSmol)
3DMet
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.029.057 Edit this at Wikidata
EC Number
  • 231-963-7
983
KEGG
MeSH Nitrous+acid
UNII
  • InChI=1S/HNO2/c2-1-3/h(H,2,3) checkY
    Key: IOVCWXUNBOPUCH-UHFFFAOYSA-N checkY
  • O=NO
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[2]
Conjugate base Nitrite
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
4
0
2
OX
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 checkY☒N ?)

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.[3] It was discovered by Carl Wilhelm Scheele, who called it "phlogisticated acid of niter". Nitrous acid is used to make diazonium salts from amines. The resulting diazonium salts are reagents in azo coupling reactions to give azo dyes.

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.[3]

Preparation and decomposition

See also: Dinitrogen trioxide

Free, gaseous nitrous acid is unstable, rapidly disproportionating to nitric oxides:

2 HNO2 → NO2 + NO + H2O

In aqueous solution, the nitrogen dioxide also disproportionates, for a net reaction producing nitric oxide and nitric acid:[4]: 1 [5]

3 HNO2 → 2 NO + HNO3 + H2O

Consequently applications of nitrous acid usually begin with mineral acid acidification of sodium nitrite. The acidification is usually conducted at ice temperatures, and the HNO2 consumed in situ.[6][7]

Nitrous acid equilibrates with dinitrogen trioxide in water, so that concentrated solutions are visibly blue:[4]: 2 

N2O3 + H2O ⇌ 2 HNO2

Addition of dinitrogen trioxide to water is thus another preparatory technique.

Chemical applications

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

At high acidities (pH ≪ 2), nitrous acid is protonated to give water and nitrosonium cations.[4]: 2 

Reduction

With I− and Fe2+ ions, NO is formed:[8]

2 HNO2 + 2 KI + 2 H2SO4 → I2 + 2 NO + 2 H2O + 2 K2SO4
2 HNO2 + 2 FeSO4 + 2 H2SO4 → Fe2(SO4)3 + 2 NO + 2 H2O + K2SO4

With Sn2+ ions, N2O is formed:

2 HNO2 + 6 HCl + 2 SnCl2 → 2 SnCl4 + N2O + 3 H2O + 2 KCl

With SO2 gas, NH2OH is formed:

2 HNO2 + 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, both HN3 and (subsequently) N2 gas are 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.[8]

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.[9] 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:[10] 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.[11][12]

DNA damage and mutation

Treatment of Escherichia coli cells with nitrous acid causes damage to the cell's DNA including deamination of cytosine to uracil, and these damages are subject to repair by specific enzymes.[13] Also, nitrous acid causes base substitution mutations in organisms with double-stranded DNA.[14]

See also

Wikimedia Commons has media related to nitrous acid.

References

  1. ^ "Nitrous Acid".
  2. ^ Perrin, D. D., ed. (1982) [1969]. Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution. IUPAC Chemical Data (2nd ed.). Oxford: Pergamon (published 1984). Entry 156. ISBN 0-08-029214-3. LCCN 82-16524.
  3. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 462.
  4. ^ a b c Williams, D. L. H. (1988). Nitrosation. Cambridge, UK: Cambridge University. ISBN 0-521-26796-X.
  5. ^ 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.
  6. ^ Petit, Y.; Larchevêque, M. (1998). "Ethyl Glycidate from (S)-Serine: Ethyl (R)-(+)-2,3-Epoxypropanoate". Org. Synth. 75: 37. doi:10.15227/orgsyn.075.0037.
  7. ^ Smith, Adam P.; Savage, Scott A.; Love, J. Christopher; Fraser, Cassandra L. (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.
  8. ^ a b Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 15: The group 15 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 449. ISBN 978-0-13-175553-6.
  9. ^ Clarke, H. T.; Kirner, W. R. (1922). "Methyl Red". Organic Syntheses. 2: 47. doi:10.15227/orgsyn.002.0047.
  10. ^ 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.
  11. ^ 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. Bibcode:2014JAWMA..64.1232S. doi:10.1080/10962247.2014.952846. PMID 25509545.
  12. ^ 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.
  13. ^ Da Roza, R.; Friedberg, E. C.; Duncan, B. K.; Warner, H. R. (1977-11-01). "Repair of nitrous acid damage to DNA in Escherichia coli". Biochemistry. 16 (22): 4934–4939. doi:10.1021/bi00641a030. ISSN 0006-2960. PMID 334252.
  14. ^ Hartman, Z.; Henrikson, E. N.; Hartman, P. E.; Cebula, T. A. (1994). "Molecular models that may account for nitrous acid mutagenesis in organisms containing double-stranded DNA". Environmental and Molecular Mutagenesis. 24 (3): 168–175. Bibcode:1994EnvMM..24..168H. doi:10.1002/em.2850240305. ISSN 0893-6692. PMID 7957120.