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Nitrite

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Nitrite
A nitrogen atom is bonded to two oxygen atoms, with bond strength 1.5, in a bent geometry; the collective ion bears a single negative charge
Space-filling model of the nitrite ion
Names
IUPAC name
Nitrite
Systematic IUPAC name
dioxidonitrate(1−)
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 233-272-6
UNII
  • InChI=1S/HNO2/c2-1-3/h(H,2,3)/p-1
    Key: IOVCWXUNBOPUCH-UHFFFAOYSA-M
  • InChI=1/HNO2/c2-1-3/h(H,2,3)/p-1
    Key: IOVCWXUNBOPUCH-REWHXWOFAR
  • N(=O)[O-]
Properties
NO−2
Molar mass 46.005 g·mol−1
Conjugate acid Nitrous acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The nitrite ion has the chemical formula NO−2. Nitrite (mostly sodium nitrite) is widely used throughout chemical and pharmaceutical industries.[1] The nitrite anion is a pervasive intermediate in the nitrogen cycle in nature. The name nitrite also refers to organic compounds having the –ONO group, which are esters of nitrous acid.

Production

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Sodium nitrite is made industrially by passing a mixture of nitrogen oxides into aqueous sodium hydroxide or sodium carbonate solution:[1][2]

NO + NO2 + 2 NaOH → 2 NaNO2 + H2O
NO + NO2 + Na2CO3 → 2 NaNO2 + CO2

The product is purified by recrystallization. Alkali metal nitrites are thermally stable up to and beyond their melting point (441 °C (826 °F) 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

Structure

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Image
The two canonical structures of NO−2, which contribute to the resonance hybrid for the nitrite ion
Image
Dimensions of trans-HONO (from the microwave spectrum)

The nitrite ion has a symmetrical structure (C2v) symmetry), with both N–O bonds having equal length and a bond angle of about 115°.[3] 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.[4] Therefore, the nitrite ion is a Lewis base.

In the gas phase, it exists predominantly as a trans-planar molecule.[citation needed]

Reactions

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Acid-base properties

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Nitrite is the conjugate base of the weak acid nitrous acid:

HNO2 ⇌ H+ + NO−2;      pKa ≈ 3.16 at 25 °C (77 °F)[5]

Nitrous acid is also highly unstable, tending to disproportionate:

3 HNO2(aq) ⇌ H3O+ + 2 NO + NO−3

This reaction is slow at 0 °C (32 °F).[2] 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

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The formal oxidation state of the nitrogen atom in nitrite is +3. This means it can be either oxidized to oxidation states +4 and +5 or reduced to as low as −3. Standard reduction potentials for reactions directly involving nitrous acid are shown in the table below:[6]

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+ → 2 Mn2+ + 3 H2O + 5 NO−3

The products of reduction reactions with the nitrite ion vary 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 (N2H+5) the product of nitrite reduction is hydrazoic acid (HN3), an unstable and explosive compound:

N2H+5 + HNO2 → 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.[2]

Analysis of nitrite

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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.[7]

Coordination complexes

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Nitrite is an ambidentate ligand and can form a wide variety of coordination complexes by binding to metal ions in several ways. For example, the red nitrito pentaamminecobalt complex [Co(NH3)5(ONO)]2+ is metastable, isomerizing to the yellow nitro complex [Co(NH3)5(NO2)]2+.[2]

Nitrite is processed by several enzymes, all of which utilize coordination complexes.[citation needed]

Hazardous reactions

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When heated with cyanides or thiosulfates, nitrites violently explode.[8]

Biochemistry

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Image
A schematic representation of the microbial nitrogen cycle.[9][10] ANAMMOX is anaerobic ammonium oxidation, DNRA is dissimilatory nitrate reduction to ammonium, and COMMAMOX is complete ammonium oxidation.

In nitrification, ammonium is converted to nitrite. Important species include Nitrosomonas. Other bacterial species, such as Nitrobacter, are responsible for oxidizing nitrite to nitrate.[citation needed]

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.[11] 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.[citation needed]

Uses

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Chemical precursor

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Azo dyes and other colorants are prepared by the process called diazotization, which requires nitrite.[1]

Meat processing and cancer risk

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The addition of nitrites and nitrates to processed meats, such as ham, bacon, and sausages, enhances the curing of meat, improves texture, imparts an attractive colour, inhibits growth of microbes and toxins, and extends shelf-life.[12][13] Nitrite reacts with meat myoglobin by attaching to the heme iron atom during cooking, forming reddish-brown nitrosomyoglobin and the characteristic pink "fresh" color of N-nitroso compounds, which can damage colon cells, potentially leading to bowel cancer.[14][15][16] Smoking, curing, salting, fermenting or adding preservatives to meats can initiate the formation of cancer-causing substances.[14][15][16][17]

Cured-meat products may be manufactured without nitrate or nitrite, such as Parma ham, which was reported in 2018 as a safe product without toxin contamination.[18] Other manufacturing processes do not assure reduction of nitrite results in toxin production.[19]

Antidote for cyanide poisoning

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Nitrites in the form of sodium nitrite and amyl nitrite are components of many cyanide antidote kits.[20] Both of these compounds bind to hemoglobin and oxidize the Fe2+ ions to Fe3+ ions forming methemoglobin. Methemoglobin, in turn, binds to cyanide (CN), creating cyanmethemoglobin, effectively removing cyanide from the complex IV of the electron transport chain (ETC) in mitochondria, which is the primary site of disruption caused by cyanide. Another mechanism by which nitrites help treat cyanide toxicity is the generation of nitric oxide, which displaces the CN from the cytochrome c oxidase (ETC complex IV), making it available for methemoglobin to bind.[21]

Organic nitrites

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Image
A nitrite ester

In organic chemistry, alkyl 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 have a vasodilating action and must be handled in the laboratory with caution. They are sometimes used in medicine to treat heart disease. 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 of nitroalkanes and nitrites.[22][23]

Safety

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Large doses of nitrites cause acute poisoning in the form of methemoglobinemia, which can lead to death.[24]

See also

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References

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  1. 1 2 3 Laue, Wolfgang; Thiemann, Michael; Scheibler, Erich; Wiegand, Karl Wilhelm (2006). "Nitrates and Nitrites". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a17_265. ISBN 978-3-527-30673-2.
  2. 1 2 3 4 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 461–464. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
  3. ↑ Goddard, John D.; Klein, Michael L. (1 August 1983). "Structure of the nitrite ion". Physical Review A. 28 (2): 1141–1143. doi:10.1103/PhysRevA.28.1141.
  4. ↑ "Molecular Orbital Theory", General Chemistry for Organic and Biological Chemistry, archived from the original on 9 January 2026, retrieved 3 July 2026
  5. ↑ da Silva, Gabriel; Kennedy, Eric M.; Dlugogorski, Bogdan Z. (1 October 2006). "Ab Initio Procedure for Aqueous-Phase pKa Calculation:  The Acidity of Nitrous Acid". The Journal of Physical Chemistry A. 110 (39): 11371–11376. doi:10.1021/jp0639243. ISSN 1089-5639.
  6. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 431. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
  7. ↑ Ivanov, V. M. (1 October 2004). "The 125th Anniversary of the Griess Reagent". Journal of Analytical Chemistry. 59 (10): 1002–1005. doi:10.1023/B:JANC.0000043920.77446.d7. ISSN 1608-3199. S2CID 98768756.
  8. ↑ Kaye, Seymour M. (1 January 1978). "N – Nitrites". Encyclopedia of Explosives and Related Items (PDF) (Technical report). Vol. 8, M1 Thickener through Pyruvonitrolic Acid. Dover, NJ: Army Armament Research And Development Center – Large Caliber Weapon Systems Lab. p. N107. LCCN 61-61759. ADA057762, PATR 2700.
  9. ↑ Sparacino-Watkins, Courtney; Stolz, John F.; Basu, Partha (16 December 2013). "Nitrate and periplasmic nitrate reductases". Chem. Soc. Rev. 43 (2): 676–706. doi:10.1039/c3cs60249d. ISSN 1460-4744. PMC 4080430. PMID 24141308.
  10. ↑ Simon, Jörg; Klotz, Martin G. (2013). "Diversity and evolution of bioenergetic systems involved in microbial nitrogen compound transformations". Biochimica et Biophysica Acta (BBA) - Bioenergetics. 1827 (2): 114–135. doi:10.1016/j.bbabio.2012.07.005. PMID 22842521.
  11. ↑ Witek J, Lakhkar AD (28 August 2023). "Nitric oxide". StatPearls, US National Library of Medicine. Retrieved 28 June 2026.
  12. ↑ "Preventive control recommendations on the use of nitrites in the curing of meat products". Canadian Food Inspection Agency, Government of Canada. 6 October 2021. Retrieved 28 June 2026.
  13. ↑ Lee, Soomin; Lee, Heeyoung; Kim, Sejeong; Lee, Jeeyeon; Ha, Jimyeong; Choi, Yukyung; Oh, Hyemin; Choi, Kyoung-Hee; Yoon, Yohan (August 2018). "Microbiological safety of processed meat products formulated with low nitrite concentration — A review". Asian-Australasian Journal of Animal Sciences. 31 (8): 1073–1077. doi:10.5713/ajas.17.0675. ISSN 1011-2367. PMC 6043430. PMID 29531192.
  14. 1 2 "Limit red and processed meat". Canadian Cancer Society. 2026. Retrieved 28 June 2026.
  15. 1 2 "Processed Meat (Sausages, Ham, Bacon, Hot Dogs, Salami)". American Institute for Cancer Research. 2025. Retrieved 28 June 2026.
  16. 1 2 "Red and Processed Meat and Cancer". American Cancer Society. 2026. Retrieved 28 June 2026.
  17. ↑ "Does processed and red meat cause cancer?". Cancer Research UK. 15 April 2025. Retrieved 28 June 2026.
  18. ↑ Wilson, Bee (1 March 2018). "Yes, bacon really is killing us". The Guardian. ISSN 0261-3077. Retrieved 14 February 2021.
  19. ↑ Lebrun, S.; Van Nieuwenhuysen, T.; Crèvecoeur, S.; Vanleyssem, R.; Thimister, J.; Denayer, S.; Jeuge, S.; Daube, G.; Clinquart, A.; Fremaux, B. (December 2020). "Influence of reduced levels or suppression of sodium nitrite on the outgrowth and toxinogenesis of psychrotrophic Clostridium botulinum Group II type B in cooked ham". International Journal of Food Microbiology. 334 108853. doi:10.1016/j.ijfoodmicro.2020.108853. PMID 32932195.
  20. ↑ Meillier, Andrew; Heller, Cara (2015). "Acute Cyanide Poisoning: Hydroxocobalamin and Sodium Thiosulfate Treatments with Two Outcomes following One Exposure Event". Case Reports in Medicine. 2015 217951. doi:10.1155/2015/217951. ISSN 1687-9627. PMC 4620268. PMID 26543483.
  21. ↑ Bebarta, Vikhyat S.; Brittain, Matthew; Chan, Adriano; Garrett, Norma; Yoon, David; Burney, Tanya; Mukai, David; Babin, Michael; Pilz, Renate B.; Mahon, Sari B.; Brenner, Matthew (June 2017). "Sodium Nitrite and Sodium Thiosulfate Are Effective Against Acute Cyanide Poisoning when Administered by Intramuscular Injection". Annals of Emergency Medicine. 69 (6): 718–725.e4. doi:10.1016/j.annemergmed.2016.09.034. ISSN 0196-0644. PMC 5446299. PMID 28041825.
  22. ↑ Victor Meyer (1872). "Über die Nitroverbindungen der Fettreihe". Justus Liebig's Annalen der Chemie (in German). 171 (1): 1–56. doi:10.1002/jlac.18741710102.
  23. ↑ Reynolds, R.B.; Adkins, H. (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. Bibcode:1929JAChS..51..279R. doi:10.1021/ja01376a037.
  24. ↑ Katabami K, Hayakawa M, Gando S (3 August 2016). "Severe Methemoglobinemia due to Sodium Nitrite Poisoning". Case Reports in Emergency Medicine. 2016 9013816. doi:10.1155/2016/9013816. PMC 4987464. PMID 27563472.
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