Nitrite
A nitrite is either a salt or an ester of nitrous acid. The nitrite ion is NO2−. The anion is bent, having the same electron distribution as ozone (O3).
Inorganic nitrites
In inorganic chemistry, nitrites are salts of nitrous acid (HNO2). They contain the nitrite ion NO2−.
Nitrites of the alkali and alkaline earth metals can be synthesized by reacting a mixture of nitrogen monoxide (NO) and nitrogen dioxide (NO2) with the corresponding metal hydroxide solution, or by the thermal decomposition of the corresponding nitrate. Other nitrites are available through reduction of the corresponding nitrates.
Nitrite is detected and analyzed by the Griess Reaction, involving the formation of a deeply red-color azo dye upon treatment of a NO2−-containing sample with sulfanilic acid and naphthyl-1-amine in the presence of acid.[1]
Nitrite can be reduced to nitric oxide or ammonia by many species of bacteria.
Organic nitrites
In organic chemistry, nitrites are esters of nitrous acid and contain the nitrosooxy functional group. They possess the general formula RONO, where R is an aryl or alkyl group. Amyl nitrite is used in medicine for the treatment of heart diseases.
Nitrites should be confused neither with nitrates, the salts of nitric acid, nor with nitro compounds, though they share the formula RNO2. The nitrite anion NO2− should not be confused with the nitronium cation NO2+.
Some specific nitrites
Sodium nitrite
Sodium nitrite (NaNO2) is a white to slight yellowish crystalline powder. It is very soluble in water and is slowly oxidized by oxygen in the air to sodium nitrate (NaNO3). It is used as a color fixative and preservative in meats and fish.
This compound is used for the curing of meat because it prevents bacterial growth and, in a reaction with the meat's myoglobin, gives the product a desirable dark red color. This nitrite, however, can be toxic at high concentrations—the lethal dose of nitrite for humans is about 22 milligrams per kilogram body weight. For this reason, the maximum allowed nitrite concentration in meat products is 200 parts per million (ppm). Under certain conditions, especially during cooking, nitrites in meat can react with degradation products of amino acids, forming nitrosamines, which are known carcinogens.
Sodium nitrite is also used in manufacturing diazo dyes, nitroso compounds, and other organic compounds; in dyeing and printing textile fabrics and bleaching fibers; in photography; as a laboratory reagent and corrosion inhibitor; in metal coatings for phosphatizing and detinning; and in the manufacture of rubber chemicals. Sodium nitrite has also been used in human and veterinary medicine as a vasodilator, a bronchodilator, an intestinal relaxant or laxative, and an antidote for cyanide poisoning.
Alkyl nitrites
Alkyl nitrites are chemical compounds with the general structure R-ONO. Formally, they are alkyl esters of nitrous acid. Methyl nitrite and ethyl nitrite are gases at room temperature and pressure, and the next few (larger) members of the series are volatile liquids. The compounds have a distinctive fruity odor.
In the laboratory, solutions of alkyl nitrites in glacial acetic acid are sometimes used as mild nitrating agents. The product formed is acetyl nitrate.
Amyl nitrite:
Amyl nitrite, also called pentyl nitrite, is an alkyl nitrite that contains five carbon atoms per molecule. Typically, the term refers to the chemical compound with the formula (CH3)2CHCH2CH2ONO. Like other volatile alkyl nitrites, it has a characteristically penetrating odor and produces marked effects on the human body when its vapor is inhaled. It acts as a vasodilator (expanding blood vessels and thus lowering blood pressure) and finds applications in medicine in the treatment of heart disease such as angina. Amyl nitrite is also used to treat cyanide poisoning by inducing the formation of methemoglobin, which sequesters cyanide as nontoxic cyanomethemoglobin.[2]
See also
Notes
- ↑ The 125th Anniversary of the Griess Reagent by V. M. Ivanov in Journal of Analytical Chemistry, Vol. 59, No. 10, 2004, pp. 1002–1005. Translated from Zhurnal Analiticheskoi Khimii, Vol. 59, No. 10, 2004, pp. 1109–1112.
- ↑ Vale, J. A. (2001). Cyanide Antidotes: from Amyl Nitrite to Hydroxocobalamin - Which Antidote is Best?. Toxicology 168 (1): 37-38.
ReferencesISBN links support NWE through referral fees
- Chang, Raymond. 2006. Chemistry. 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031.
- Cotton, F. Albert, and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry. 4th ed. New York: Wiley. ISBN 0-471-02775-8.
External links
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