Nitrite

Nitrite ion with an O-N-O bond angle of ca. 120°

Space-filling model of the nitrite ion

The nitrite ion has the chemical formula NO₂⁻. The anion is symmetric with equal N-O bond lengths and a O-N-O bond angle of ca. 120°. On protonation the unstable weak acid nitrous acid is produced. Nitrite can be oxidised or reduced, with product somewhat dependent on the oxidizing/reducing agent. The nitrite ion is an ambidentate ligand and is known to bond to metal centres in at least five different ways.^[1] Nitrite is important in biochemistry as a source of the vasodilator nitric oxide. Nitrites are used for curing meat. In organic chemistry the NO₂ group is present in nitrous acid esters and nitro compounds.

The nitrite ion

Nitrite salts

Sodium nitrite is made industrially by passing nitrous fumes into aqueous sodium hydroxide or sodium carbonate solution.^[1]

NO + NO₂ + 2NaOH (or Na₂CO₃) → 2NaNO₂ +H₂O ( or CO₂)

The product is purified by recrystallization. Alkali metal nitrites are thermally stable up to and beyond the melting point (441°C for KNO₂). Ammonium nitrite can be made from dinitrogen trioxide, N₂O₃, which is formally the anhydride of nitrous acid.

2NH₃ + H₂O +N₂O₃ → 2NH₄NO₂

This compound may decompose explosively on heating.

In organic chemistry nitrites are used in diazotization reactions.

Structure

The two canonical structures of NO₂⁻, which contribute to the resonance hybrid for the nitrite ion

The nitrite ion has a symmetrical structure (C_2v symmetry) with both N-O bonds having equal length. 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 are 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.

Acid-base properties

Dimensions of trans HONO (from the microwave spectrum)

In aqueous solution nitrous acid is a weak acid.

HNO₂ H⁺ + NO₂^-; pK_a = ca. 3.3 at 18°C^[2]

Nitrous acid is volatile; in the gas phase it exists predominantly as a trans- planar molecule. In solution it is unstable with respect to the disproportionation reaction

3HNO₂ (aq) H₃O⁺ + NO₃^- + 2NO

This reaction is slow at 0°C.^[1] Addition of acid to a solution of a nitrite in the presence of a reducing agent such as iron(II) is a way to make nitric oxide, NO, in the laboratory.

Oxidation and reduction

The formal oxidation state of the nitrogen atom in a nitrite is +3. This means that it is can be either oxidised 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.^[3]

Half-reaction	E0/V
NO3- + 3H+ + 2e- HNO2 + H2O	+0.94
2HNO2+ 4H+ + 4e- H2N2O2 + 2H2O	+0.86
N2O4 + 2H+ + 2e- 2HNO2	+1.065
2HNO2+ 4H+ + 4e- N2O + 3H2O	+1.29

The data can be extended to include products in lower oxidation states. For example,

H₂N₂O₂ + 2H⁺ + 2e^- N₂ + 2H₂O; E⁰ = 2.65V

Oxidation reactions usually result in the formation of the nitrate ion, with nitrogen in oxidation state +5. For example, oxidation with permanganate can be used for quantitative analysis of nitrite, by titration.

5NO₂^- + 2MnO₄^- + 6H⁺ → 5NO₃^- + 2Mn²⁺ + 3H₂O

The product of reduction reactions are various depending on the reducing agent used. 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 occurs with hydrogen sulfide. With the hydrazinium cation, N₂H₅⁺, hydrogen azide, HN₃, is produced

HNO₂ + N₂H₅⁺ → HN₃ + H₂O + H₃O⁺

which can also further react with nitrite

HNO₂ + HN₃ → N₂O + N₂ + H₂O

This reaction is unusual in that it involves compounds with nitrogen in four different oxidation states.^[1]

Coordination complexes

The nitrite ion is known to form complexes in at least five different ways.^[1]

1. When donation is from nitrogen to a metal centre, the complex is known as a nitro- complex.
2. When donation is from one oxygen to a metal centre, the complex is known as a nitrito- complex.
3. Both oxygen atom may donate to a metal centre, forming a chelate complex.
4. A nitrite ion can form an unsymmetrical bridge between two metal centres, donating through nitrogen to one metal and oxygen to the other.
5. A single oxygen atom can bridge to two metal centres.

Alfred Werner studied the nitro-nitrito isomerism (1 and 2) extensively. The red isomer of cobalt pentammine 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 and the two bipy ligands occupy four coordination sites on the copper ion so 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.^[1]

Nitrite in biochemistry

Sodium nitrite 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. Because of the toxicity of nitrite (the lethal dose of nitrite for humans is about 22 mg per kg body weight), the maximum allowed nitrite concentration in meat products is 200 ppm. Under certain conditions, especially during cooking, nitrites in meat can react with degradation products of amino acids, forming nitrosamines, which are known carcinogens.^[4]

Nitrite is detected and analyzed by the Griess Reaction, involving the formation of a deep red-colored azo dye upon treatment of a NO₂⁻-containing sample with sulfanilic acid and naphthyl-1-amine in the presence of acid.^[5] 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.

Organic nitrites and nitro compounds

A nitrite ester

aromatic nitration

In organic chemistry, nitrites are esters of nitrous acid and contain the nitrosooxy 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 is used in medicine for the treatment of heart diseases. A classic named reaction for the synthesis of alkyl nitrites is the Meyer synthesis^[6][7] in which alkyl halides react with metallic nitrites to a mixture to nitroalkanes and nitrites.

Nitrobenzene is a simple example of a nitro compound. In aromatic nitration reaction a C-H bond is broken leaving the two electron on the carbon atom. The two electrons are added to the nitronium ion reducing it to nitrite.

See also

• Curing (food preservation)

References

1. ^ ^{a b c d e f} Greenwood, pp 461-464
2. ^ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
3. ^ Greenwood, p 431
4. ^ Jakszyn P, Gonzalez CA. Nitrosamine and related food intake and gastric and oesophageal cancer risk: a systematic review of the epidemiological evidence. World J Gastroenterol. 2006 Jul 21;12(27):4296-303. PubMed
5. ^ V. M. Ivanov (2004). "The 125th Anniversary of the Griess Reagent". Journal of Analytical Chemistry 59 (10): 1002–1005. doi:10.1023/B:JANC.0000043920.77446.d7.  Translated from V. M. Ivanov (2004). Zhurnal Analiticheskoi Khimii 59 (10): 1109–1112. 
6. ^ Victor Meyer (1872). "Ueber die Nitroverbindungen der Fettreihe". Justus Liebig's Annalen der Chemie 171 (1): 1–56. doi:10.1002/jlac.18741710102. ; Victor Meyer, J. Locher (1876). "Ueber die Pseudonitrole, die Isomeren der Nitrolsäuren". Justus Liebig's Annalen der Chemie 180 (1-2): 133–155. doi:10.1002/jlac.18761800113. ; V. Meyer and Stüber (1872). "Vorläufige Mittheilung". Chemische Berichte 5: 203–205. doi:10.1002/cber.18720050165. ; Victor Meyer, O. Stüber (1872). "Ueber die Nitroverbindungen der Fettreihe". Chemische Berichte 5: 399. doi:10.1002/cber.187200501121. ; Victor Meyer, A. Rilliet (1872). "Ueber die Nitroverbindungen der Fettreiche. Dritte Mittheilung". Chemische Berichte 5: 1029–1034. doi:10.1002/cber.187200502133. ; Victor Meyer, C. Chojnacki (1872). "Ueber die Nitroverbindungen der Fettreihe. Vierte Mittheilung". Chemische Berichte 5: 1034–1038. doi:10.1002/cber.187200502134. 
7. ^ Robert B. Reynolds, Homer Adkins (1929). "The Relationship of the Constitution of Certain Alky Halides to the Formation of Nitroparaffins and Alkyl Nitrites". Journal of the American Chemical Society 51 (1): 279–287. doi:10.1021/ja01376a037. 

External links

• Material Safety Data Sheet, sodium nitrite
• ATSDR - Case Studies in Environmental Medicine - Nitrate/Nitrite Toxicity U.S. Department of Health and Human Services (public domain)

Bibliography

• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0080379419.