Nitrite

Nitrite
Space-filling model of the nitrite ion
Names
Systematic IUPAC name
dioxidonitrate(1−)
Other names
nitrite
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number 233-272-6
Properties
NO−
2
Molar mass 46.01 g·mol−1
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

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.[1] Nitrite is also important in biochemistry as a source of the potent vasodilator nitric oxide. In organic chemistry the –NO2 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.[2][3]

The nitrite ion

Nitrite salts

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

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:

2 NH3 + H2O + N2O3 → 2 NH4NO2

This compound may decompose explosively on heating.

In organic chemistry nitrites are used in diazotization reactions.

Structure

The two canonical structures of NO−
2, which contribute to the resonance hybrid for the nitrite ion.

The nitrite ion has a symmetrical structure (C2v 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 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.

Acid-base properties

Dimensions of trans-HONO (from the microwave spectrum)

In aqueous solution, nitrous acid is a weak acid:

HNO2 ⇌ H+ + NO−
2;      pKa ≈ 3.3 at 18 °C[4]

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:

3HNO2 (aq) ⇌ H3O+ + NO−
3 + 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 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:[5]

Half-reactionE0 (V)
NO−
3 + 3 H+ + 2 e− ⇌ HNO2 + H2O		 +0.94
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:

H2N2O2 + 2 H+ + 2 e− ⇌ N2 + 2 H2O;      E0 = +2.65 V

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):

5 NO−
2 + 2 MnO−
4 + 6 H+ → 5 NO−
3 + 2 Mn2+ + 3 H2O

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
2H+
5) the product is hydrazoic acid (HN3), an explosive compound:

HNO2 + N
2H+
5 → 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 coordination complexes in at least five different ways.[1]

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

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.[1]

Nitrite in food preservation and biochemistry

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.[6] 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.[7] They are considered irreplaceable in the prevention of botulinum poisoning from consumption of cured dry sausages by preventing spore germination.[8] 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.[9]

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.[10] 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.

Organic nitrites and nitro compounds

A nitrite ester
Aromatic nitration

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[11][12] 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 reactions a C–H bond is broken, leaving the two electrons on the carbon atom. These two electrons are added to the nitronium ion, reducing it to nitrite.

See also

References

  1. 1 2 3 4 5 6 Greenwood, pp 461–464
  2. ↑ "List of classifications, Volumes 1–116 - IARC Monographs on the Evaluation of Carcinogenic Risks to Humans". International Agency for Research on Cancer (IARC) - World Health Organization (WHO). 2010. Retrieved 25 September 2016.
  3. ↑ "VOLUME 94 - Ingested Nitrate and Nitrite, and Cyanobacterial Peptide Toxins - IARC Monographs on the Evaluation of Carcinogenic Risks to Humans". International Agency for Research on Cancer (IARC) - World Health Organization (WHO). 2010. Retrieved 25 September 2016.
  4. ↑ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
  5. ↑ Greenwood, p 431
  6. ↑ Jakszyn, P.; Gonzalez, C. A. (2006). "Nitrosamine and related food intake and gastric and oesophageal cancer risk: A systematic review of the epidemiological evidence". World journal of gastroenterology: WJG. 12 (27): 4296–4303. PMC 4087738 Freely accessible. PMID 16865769. doi:10.3748/wjg.v12.i27.4296.
  7. ↑ sodium nitrite and nitrate facts Accessed Dec 12, 2014
  8. ↑ De Vries, John (1997). Food Safety and Toxicity. CRC Press. p. 70. ISBN 978-0-8493-9488-1.
  9. ↑ Charles, R. L.; Rudyk, O.; Prysyazhna, O.; Kamynina, A.; Yang, J.; Morisseau, C.; Hammock, B. D.; Freeman, B. A.; Eaton, P. (2014). "Protection from hypertension in mice by the Mediterranean diet is mediated by nitro fatty acid inhibition of soluble epoxide hydrolase". Proceedings of the National Academy of Sciences. 111 (22): 8167. PMC 4050620 Freely accessible. PMID 24843165. doi:10.1073/pnas.1402965111.
  10. ↑ 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. Missing or empty |title= (help)
  11. ↑ 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): 133–55. doi:10.1002/jlac.18761800113.; V. Meyer and Stüber (1872). "Vorläufige Mittheilung". Chemische Berichte. 5: 203–05. 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 (2): 1029–34. doi:10.1002/cber.187200502133.; Victor Meyer, C. Chojnacki (1872). "Ueber die Nitroverbindungen der Fettreihe. Vierte Mittheilung". Chemische Berichte. 5 (2): 1034–38. doi:10.1002/cber.187200502134.
  12. ↑ 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): 279–87. doi:10.1021/ja01376a037.

Bibliography

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