Difference between revisions of "Nitrous acid" - New World Encyclopedia

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Nitrous acid

General
Systematic name	Dioxonitric(III) acid
Other names	Nitrous acid
Molecular formula	H N O₂
Molar mass	47.0134 g/mol
CAS number	7782-77-6
Properties
Density	? g/cm³
Solubility (water)
Melting point	? °C
Boiling point	? °C
Acid dissociation constant pK_a	3.34
Disclaimer and references

Nitrous acid (molecular formula H N O₂) is a weak acid known only in solution and in the form of nitrite salts. In the atmosphere, it is formed as an intermediate, breaking down to produce hydroxyl free radicals. These free radicals are then involved in regulating the amount of ozone in the troposphere (lower atmosphere).

Nitrous acid is useful for destroying sodium azide solutions, which are toxic and potentially explosive. It can also be used to prepare diazonium salts, which combine with anilines and phenols to form brightly colored azo compounds. The latter type of reaction can be used to produce azo-dyes, as well as serving as a qualitative test for aromatic amines.

Preparation

Nitrous acid is readily formed in solution by the action of any mineral acid on sodium nitrite (NaNO₂). The reaction may be written as follows:

NaNO₂ + H⁺ ----> HNO₂ + Na⁺

Chemical reactions

Acid-base reactions

Nitrous acid is a monobasic acid, which means that each molecule of HNO₂ releases one proton (H⁺) in solution. In addition, it is a weak acid, which means that at any given time, the dissociated molecules are in equilibrium with undissociated molecules.

Also, like other acids, nitrous acid reacts with a base to form a salt. In this case, sodium nitrite is formed.

Decomposition reactions

Nitrous acid is unstable, and it may decompose in two ways, in solution.

In one type of decomposition, it produces nitrogen dioxide, nitric oxide, and water, as follows:

2HNO₂ ----> NO₂ + NO + H₂O

In a second pathway, it may decompose to nitric acid, nitrous oxide, and water, as follows:

4HNO₂ ----> 2HNO₃ + N₂O + H₂O

Atmospheric relevance

Nitrous acid is formed in the atmosphere as an important intermediate. It is produced by the reaction of nitrogen dioxide (NO₂) and water on various surfaces such as atmospheric aerosols. The action of sunlight then decomposes it to produce hydroxyl free radicals, which are intricately involved in regulating the ozone budget of the troposphere (lower atmosphere).

Applications

Nitrous acid is used to destroy toxic and potentially explosive sodium azide solutions. Nitrous acid can be used to prepare diazonium salts, which couple with anilines and phenols to form brightly colored azo compounds in a qualitative test for aromatic amines. This reaction is also used to produce azo-dyes.

References
ISBN links support NWE through referral fees

Chang, Raymond. 2006. Chemistry. 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031 and ISBN 978-0073221038.

Cotton, F. Albert, Geoffrey Wilkinson, Carlos A. Murillo, and Manfred Bochmann. 1999. Advanced Inorganic Chemistry. 6th edition. New York: Wiley. ISBN 0471199575

Greenwood, N.N., and A. Earnshaw. 1998. Chemistry of the Elements. 2nd ed. Oxford, UK; Burlington, MA: Butterworth-Heinemann, Elsevier Science. ISBN 0750633654. Online version.

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@@ Line 47: / Line 47: @@
 |-
 |}
-'''Nitrous acid''' (molecular formula [[Hydrogen|H]][[Nitrogen|N]][[Oxygen|O]]<sub>2</sub>) is a weak 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 as a [[carbene]].
+'''Nitrous acid''' (molecular formula [[Hydrogen|H]][[Nitrogen|N]][[Oxygen|O]]<sub>2</sub>) is a weak [[acid]] known only in [[solution]] and in the form of [[nitrite]] salts. In the atmosphere, it is formed as an intermediate, breaking down to produce hydroxyl [[free radical]]s. These free radicals are then involved in regulating the amount of [[ozone]] in the [[troposphere]] (lower atmosphere).
+Nitrous acid is useful for destroying [[sodium azide]] solutions, which are toxic and potentially explosive. It can also be used to prepare [[diazonium salt]]s, which combine with [[aniline]]s and [[phenol]]s to form brightly colored [[azo compound]]s. The latter type of reaction can be used to produce [[azo-dye]]s, as well as serving as a qualitative test for aromatic amines.
 ==Preparation==
-It can be prepared by adding any mineral acid to [[sodium nitrite]].
-==Decomposition==
+Nitrous acid is readily formed in solution by the action of any mineral acid on [[sodium nitrite]] (NaNO<sub>2</sub>). The reaction may be written as follows:
-It rapidly decomposes into [[nitrogen dioxide]], [[nitric oxide]] and [[water]] when in solution.
+::NaNO<sub>2</sub> + H<sup>+</sup>  ---->  HNO<sub>2</sub> + Na<sup>+</sup>
+== Chemical reactions ==
+=== Acid-base reactions ===
+Nitrous acid is a monobasic acid, which means that each molecule of HNO<sub>2</sub> releases one proton (H<sup>+</sup>) in solution. In addition, it is a weak acid, which means that at any given time, the dissociated molecules are in equilibrium with undissociated molecules.
+Also, like other acids, nitrous acid reacts with a base to form a salt. In this case, sodium nitrite is formed.
-HNO<sub>2</sub>------>NO<sub>2</sub> + NO + H<sub>2</sub>O
+===Decomposition reactions===
-It also decomposes into [[nitric acid]] and [[nitrous oxide]] and water
+Nitrous acid is unstable, and it may decompose in two ways, in solution.
-HNO<sub>2</sub> -> 2HNO<sub>3</sub> + N<sub>2</sub>O + H<sub>2</sub>O
+In one type of decomposition, it produces [[nitrogen dioxide]], [[nitric oxide]], and [[water]], as follows:
+::2HNO<sub>2</sub>  ---->  NO<sub>2</sub> + NO + H<sub>2</sub>O
-==Chemistry==
+In a second pathway, it may decompose to [[nitric acid]], [[nitrous oxide]], and water, as follows:
-Nitrous acid can be used to prepare [[diazonium salt]]s which couple with anilines and phenols to form brightly colored [[azo compound]]s in a qualitative test for aromatic amines. This reaction is also used to produce [[azo-dyes]]. Nitrous acid is used to destroy toxic and potentially explosive [[sodium azide]] solutions. Nitrous acid is usually formed ''in situ'' by the action of mineral acid on [[sodium nitrite]].
+::4HNO<sub>2</sub>  ---->  2HNO<sub>3</sub> + N<sub>2</sub>O + H<sub>2</sub>O
 ==Atmospheric relevance==
-Nitrous acid is an important atmospheric intermediate. It is produced by the heterogeneous reaction of NO<sub>2</sub> and water on various surfaces such as atmospheric aerosols. It is readily photolysed to produce hydroxyl radicals which are intricately involved in the ozone budget of the troposphere (lower atmosphere).
+Nitrous acid is formed in the atmosphere as an important intermediate. It is produced by the reaction of nitrogen dioxide (NO<sub>2</sub>) and water on various surfaces such as atmospheric aerosols. The action of sunlight then decomposes it to produce hydroxyl [[free radical]]s, which are intricately involved in regulating the [[ozone]] budget of the troposphere (lower atmosphere).
+== Applications ==
+Nitrous acid is used to destroy toxic and potentially explosive [[sodium azide]] solutions. Nitrous acid can be used to prepare [[diazonium salt]]s, which couple with [[aniline]]s and [[phenol]]s to form brightly colored [[azo compound]]s in a qualitative test for aromatic amines. This reaction is also used to produce [[azo-dye]]s.
 ==See also==
-* [[Demjanov rearrangement]]
-* [[Nitric acid]] ([[Hydrogen|H]][[Nitrogen|N]][[Oxygen|O]]<sub>3</sub>)
-* [[Tiffeneau-Demjanov rearrangement]]
-==External links==
+* [[Acid]]
+* [[Nitric acid]]
+* [[Sulfuric acid]]
+== References ==
+* Chang, Raymond. 2006. ''Chemistry''. 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031 and ISBN 978-0073221038.
+* Cotton, F. Albert, Geoffrey Wilkinson, Carlos A. Murillo, and Manfred Bochmann. 1999. ''Advanced Inorganic Chemistry''. 6th edition. New York: Wiley. ISBN 0471199575
-{{ChemicalSources}}
+* Greenwood, N.N., and A. Earnshaw. 1998. ''Chemistry of the Elements''. 2nd ed. Oxford, UK; Burlington, MA: Butterworth-Heinemann, Elsevier Science. ISBN 0750633654. [http://www.knovel.com/knovel2/Toc.jsp?BookID=402&VerticalID=0 Online version].
 [[Category:Physical sciences]]