Nitrous acid
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| Names | |
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Preferred IUPAC name Nitrous acid | |
Systematic IUPAC name Hydroxidooxidonitrogen | |
| Identifiers | |
CAS Number |
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3D model (JSmol) |
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3DMet | B00022 |
ChEBI |
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ChEMBL |
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ChemSpider |
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ECHA InfoCard | 100.029.057 |
EC Number | 231-963-7 |
Gmelin Reference | 983 |
KEGG |
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MeSH | Nitrous+acid |
PubChem CID |
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InChI
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SMILES
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| Properties | |
Chemical 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 |
Conjugate base | Nitrite |
| Hazards | |
NFPA 704 |  0 4 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  Y N ?) | |
Infobox references | |
Nitrous acid (molecular formula HNO2) is a weak and monobasic acid known only in solution and in the form of nitrite 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.
Contents
1 Structure
2 Preparation
3 Reactions
3.1 Decomposition
3.2 Reduction
3.3 Organic chemistry
4 Atmosphere of the Earth
5 See also
6 References
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
Nitrous acid is usually generated by acidification of aqueous solutions 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 oxide in water according to the equation
- N2O3 + H2O → 2 HNO2
Reactions
Decomposition
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 trioxide 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
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
 | Wikimedia Commons has media related to nitrous acid. |
- Demjanov rearrangement
Nitric acid (HNO3)- Tiffeneau-Demjanov rearrangement
References
^ ab Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9..mw-parser-output cite.citation{font-style:inherit}.mw-parser-output q{quotes:"""""""'""'"}.mw-parser-output code.cs1-code{color:inherit;background:inherit;border:inherit;padding:inherit}.mw-parser-output .cs1-lock-free a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/6/65/Lock-green.svg/9px-Lock-green.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-lock-limited a,.mw-parser-output .cs1-lock-registration a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/d/d6/Lock-gray-alt-2.svg/9px-Lock-gray-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-lock-subscription a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/a/aa/Lock-red-alt-2.svg/9px-Lock-red-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration{color:#555}.mw-parser-output .cs1-subscription span,.mw-parser-output .cs1-registration span{border-bottom:1px dotted;cursor:help}.mw-parser-output .cs1-hidden-error{display:none;font-size:100%}.mw-parser-output .cs1-visible-error{font-size:100%}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration,.mw-parser-output .cs1-format{font-size:95%}.mw-parser-output .cs1-kern-left,.mw-parser-output .cs1-kern-wl-left{padding-left:0.2em}.mw-parser-output .cs1-kern-right,.mw-parser-output .cs1-kern-wl-right{padding-right:0.2em} p. 462
^ 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)
^ 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)
^
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.
^ ab 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.
^ 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)
^ Prudent practices in the laboratory: handling and disposal of chemicals. Washington, D.C.: National Academy Press. 1995. ISBN 978-0-309-05229-0.
^ 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. PMID 25509545.
^ 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. doi:10.1021/acs.jpca.7b10715. PMID 29182863.
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