The nitrite ion, which has the chemical formula NO−
2, is a symmetric anion with equal N–O bond lengths. Upon protonation, the unstable weak acid nitrous acid is produced. Nitrite can be oxidized or reduced, with the product somewhat dependent on the oxidizing/reducing agent and its strength. The nitrite ion is an ambidentate ligand, and is known to bond to metal centers in at least five different ways. Nitrite is also important in biochemistry as a source of the potent vasodilator nitric oxide. In organic chemistry the NO−
2 group is present in nitrous acid esters and nitro compounds. Nitrites are also used in the food production industry for curing meat.
Nitrate or nitrite (ingested) under conditions that result in endogenous nitrosation has been classified as "Probably carcinogenic to humans" (Group 2A) by International Agency for Research on Cancer (IARC), the specialized cancer agency of the World Health Organization (WHO) of the United Nations.
The product is purified by recrystallization. Alkali metal nitrites are thermally stable up to and beyond their melting point (441 °C for KNO2). Ammonium nitrite can be made from dinitrogen trioxide, N2O3, which is formally the anhydride of nitrous acid:
This compound may decompose explosively on heating.
In organic chemistry nitrites are used in diazotization reactions.
The nitrite ion has a symmetrical structure (C2v symmetry), with both N–O bonds having equal length and a bond angle of about 115 deg. In valence bond theory, it is described as a resonance hybrid with equal contributions from two canonical forms that are mirror images of each other. In molecular orbital theory, there is a sigma bond between each oxygen atom and the nitrogen atom, and a delocalized pi bond made from the p orbitals on nitrogen and oxygen atoms which is perpendicular to the plane of the molecule. The negative charge of the ion is equally distributed on the two oxygen atoms. Both nitrogen and oxygen atoms carry a lone pair of electrons. Therefore, the nitrite ion is a Lewis base. Moreover, it can act as an ambidentate ligand towards a metal ion, donating a pair of electrons from either nitrogen or oxygen atoms.
Nitrous acid is also highly volatile – in the gas phase it exists predominantly as a trans-planar molecule. In solution, it is unstable with respect to the disproportionation reaction:
The formal oxidation state of the nitrogen atom in a nitrite is +3. This means that it can be either oxidized to oxidation states +4 and +5, or reduced to oxidation states as low as −3. Standard reduction potentials for reactions directly involving nitrous acid are shown in the table below:
3 + 3 H+ + 2 e− ⇌ HNO2 + H2O
|2 HNO2 + 4 H+ + 4 e− ⇌ H2N2O2 + 2 H2O||+0.86|
|N2O4 + 2 H+ + 2 e− ⇌ 2 HNO2||+1.065|
|2 HNO2+ 4 H+ + 4 e− ⇌ N2O + 3 H2O||+1.29|
The data can be extended to include products in lower oxidation states. For example:
Oxidation reactions usually result in the formation of the nitrate ion, with nitrogen in oxidation state +5. For example, oxidation with permanganate ion can be used for quantitative analysis of nitrite (by titration):
The product of reduction reactions with nitrite ion are varied, depending on the reducing agent used and its strength. With sulfur dioxide, the products are NO and N2O; with tin(II) (Sn2+) the product is hyponitrous acid (H2N2O2); reduction all the way to ammonia (NH3) occurs with hydrogen sulfide. With the hydrazinium cation (N
5) the product is hydrazoic acid (HN3), an explosive compound:
which can also further react with nitrite:
This reaction is unusual in that it involves compounds with nitrogen in four different oxidation states.
The nitrite ion is known to form coordination complexes in at least five different ways.
Alfred Werner studied the nitro–nitrito isomerism (1 and 2) extensively. The red isomer of pentaamminecobalt(III) with nitrite is now known to be a nitrito complex, [Co(NH3)5(ONO)]2+; it is metastable and isomerizes to the yellow nitro complex [Co(NH3)5(NO2)]2+. An example of chelating nitrite (3) was found in [Cu(bipy)2(O2N)]NO3 – "bipy" is the bidentate ligand 2,2′-bypyridyl, with the two bipy ligands occupying four coordination sites on the copper ion, so that the nitrite is forced to occupy two sites in order to achieve an octahedral environment around the copper ion. Examples of 4 and 5 are illustrated.
Sodium nitrite is used for the curing of meat because it prevents bacterial growth and, as it is a reducing agent (opposite of oxidation agent), in a reaction with the meat's myoglobin, gives the product a desirable pink-red "fresh" color, such as with corned beef. This use of nitrite goes back to the Middle Ages, and in the US has been formally used since 1925. Because of the relatively high toxicity of nitrite (the lethal dose in humans is about 22 milligrams per kilogram of body weight), the maximum allowed nitrite concentration in meat products is 200 ppm. At these levels, some 80 to 90% of the nitrite in the average U.S. diet is not from cured meat products, but from natural nitrite production from vegetable nitrate intake. Under certain conditions – especially during cooking – nitrites in meat can react with degradation products of amino acids, forming nitrosamines, which are known carcinogens. However, the role of nitrites (and to some extent nitrates) in preventing botulism by preventing C. botulinum endospores from germinating have prevented the complete removal of nitrites from cured meat, and indeed by definition in the U.S., meat cannot be labeled as "cured" without nitrite addition. They are considered irreplaceable in the prevention of botulinum poisoning from consumption of cured dry sausages by preventing spore germination. To reduce nitrosamine generation, sodium ascorbate or its stereoisomer sodium erythorbate may be added to cured meat. In mice, food rich in nitrites together with unsaturated fats can prevent hypertension, which is one explanation for the apparent health effect of the Mediterranean diet.
Nitrite is detected and analyzed by the Griess Reaction, involving the formation of a deep red-colored azo dye upon treatment of a NO−
2-containing sample with sulfanilic acid and naphthyl-1-amine in the presence of acid. Nitrite can be reduced to nitric oxide or ammonia by many species of bacteria. Under hypoxic conditions, nitrite may release nitric oxide, which causes potent vasodilation. Several mechanisms for nitrite conversion to NO have been described, including enzymatic reduction by xanthine oxidoreductase, nitrite reductase, and NO synthase (NOS), as well as nonenzymatic acidic disproportionation reactions.
In organic chemistry, nitrites are esters of nitrous acid and contain the nitrosoxy functional group. Nitro compounds contain the C–NO2 group. Nitrites have the general formula RONO, where R is an aryl or alkyl group. Amyl nitrite and other alkyl nitrites are used in medicine for the treatment of heart diseases, and occasionally used recreationally for their "rush" and prolongation of orgasm, particularly in males. A classic named reaction for the synthesis of alkyl nitrites is the Meyer synthesis in which alkyl halides react with metallic nitrites to a mixture to nitroalkanes and nitrites.
For example, the USDA now requires adding ascorbic acid (vitamin C) or erythorbic acid to bacon cure, a practice that greatly reduces the formation of nitrosamines.
|Wikimedia Commons has media related to Nitrites.|