| Nitrous acid |
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| 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 Y |
| Gmelin Reference |
983 |
| 3DMet |
B00022 |
| Jmol-3D images |
Image 1 |
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InChI=1S/HNO2/c2-1-3/h(H,2,3) Y
Key: IOVCWXUNBOPUCH-UHFFFAOYSA-N Y
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| 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
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| 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.
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Contents
- 1 Structure
- 2 Preparation
- 3 Decomposition
- 4 Chemistry
- 4.1 Inorganic Chemistry
- 4.2 Organic Chemistry
- 5 Atmosphere of the earth
- 6 See also
- 7 References
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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]
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dimensions of the trans form
(from the microwave spectrum)
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model of the trans form
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cis form
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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
Inorganic Chemistry
Nitrous acid can be oxidized to nitrate by powerful oxidizing agents, such as acidified potassium permanganate:[3]
- 3 H2SO4 + 2 KMnO4 + 5 KNO2 → 5 KNO3 + 3 H2O + MnSO4 + K2SO4
Reduction of the acid gives different products, depending on the reducing agent:[3]
With I- and Fe2+ ions, NO is formed:
- 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 + 2 N2O + 6 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 N2H5+, 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}
- NO3- + 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 Ecello 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 more faster, rather than a more powerful oxidizing agent than dilute nitric aid.[3]
Organic 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.[4] 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:[5] It is aminly 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)2CH-CH2-CH2-OH + HNO2 → (CH3)2CH-CH2-CH2-ONO + H2O
Atmosphere of the earth
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This section does not cite any references or sources. (June 2010) |
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.
See also
- Demjanov rearrangement
- Nitric acid (HNO3)
- Tiffeneau-Demjanov rearrangement
References
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth–Heinemann. ISBN 0080379419. p. 462
- ^ 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.
- ^ a b c 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). http://www.orgsyn.org/orgsyn/pdfs/CV1P0374.pdf
- ^ Prudent practices in the laboratory: handling and disposal of chemicals. Washington, D.C.: National Academy Press. 1995. ISBN 0-309-05229-7. http://books.nap.edu/openbook.php?record_id=4911&page=165.
