Nitric oxide (nitrogen oxide or nitrogen monoxide[1]) is a colorless gas with the formula NO. It is one of the principal oxides of nitrogen. Nitric oxide is a free radical: it has an unpaired electron, which is sometimes denoted by a dot in its chemical formula (N=O or NO). Nitric oxide is also a heteronuclear diatomic molecule, a class of molecules whose study spawned early modern theories of chemical bonding.[6]

Nitric oxide
Skeletal formula of nitric oxide with bond length
Skeletal formula showing two lone pairs and one three-electron bond
Skeletal formula showing two lone pairs and one three-electron bond
Space-filling model of nitric oxide
Space-filling model of nitric oxide
IUPAC name
Nitrogen monoxide[1]
Systematic IUPAC name
Oxidonitrogen(•)[2] (additive)
Other names
Nitrogen oxide
Nitrogen(II) oxide
Nitrogen monoxide
3D model (JSmol)
ECHA InfoCard 100.030.233 Edit this at Wikidata
EC Number
  • 233-271-0
RTECS number
  • QX0525000
UN number 1660
  • InChI=1S/NO/c1-2 checkY
  • InChI=1/NO/c1-2
  • [N]=O
Molar mass 30.006 g·mol−1
Appearance Colourless gas
Density 1.3402 g/L
Melting point −164 °C (−263 °F; 109 K)
Boiling point −152 °C (−242 °F; 121 K)
0.0098 g / 100 ml (0 °C)
0.0056 g / 100 ml (20 °C)
linear (point group Cv)
210.76 J/(K·mol)
90.29 kJ/mol
R07AX01 (WHO)
License data
via pulmonary capillary bed
2–6 seconds
Occupational safety and health (OHS/OSH):
Main hazards
  • Fatal if inhaled
  • Causes severe burns
  • Causes eye damage
  • Corrosive to the respiratory tract
GHS labelling:
GHS04: Compressed GasGHS03: OxidizingGHS05: CorrosiveGHS06: Toxic[3][4]
H270, H280, H314, H330[3][4]
P220, P244, P260, P280, P303+P361+P353+P315, P304+P340+P315, P305+P351+P338+P315, P370+P376, P403, P405[3][4]
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
315 ppm (rabbit, 15 min)
854 ppm (rat, 4 h)
2500 ppm (mouse, 12 min)[5]
320 ppm (mouse)[5]
Safety data sheet (SDS) External SDS
Related compounds
Related nitrogen oxides
Dinitrogen pentoxide

Dinitrogen tetroxide
Dinitrogen trioxide
Nitrogen dioxide
Nitrous oxide
Nitroxyl (reduced form)
Hydroxylamine (hydrogenated form)

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

An important intermediate in industrial chemistry, nitric oxide forms in combustion systems and can be generated by lightning in thunderstorms. In mammals, including humans, nitric oxide is a signaling molecule in many physiological and pathological processes.[7] It was proclaimed the "Molecule of the Year" in 1992.[8] The 1998 Nobel Prize in Physiology or Medicine was awarded for discovering nitric oxide's role as a cardiovascular signalling molecule.[9]

Nitric oxide should not be confused with nitrogen dioxide (NO2), a brown gas and major air pollutant, or with nitrous oxide (N2O), an anesthetic gas.[6]

Physical properties


Electronic configuration


The ground state electronic configuration of NO is, in united atom notation:[10]


The first two orbitals are actually pure atomic 1sO and 1sN from oxygen and nitrogen respectively and therefore are usually not noted in the united atom notation. Orbitals noted with an asterisk are antibonding. The ordering of 5σ and 1π according to their binding energies is subject to discussion. Removal of a 1π electron leads to 6 states whose energies span over a range starting at a lower level than a 5σ electron an extending to a higher level. This is due to the different orbital momentum couplings between a 1π and a 2π electron.

The lone electron in the 2π orbital makes NO a doublet (X ²Π) in its ground state whose degeneracy is split in the fine structure from spin-orbit coupling with a total momentum J=32 or J=12.



The dipole of NO has been measured experimentally to 0.15740 D and is oriented from O to N (⁻NO⁺) due to the transfer of negative electronic charge from oxygen to nitrogen. [11]



With di- and triatomic molecules


Upon condensing to a liquid, nitric oxide dimerizes to dinitrogen dioxide, but the association is weak and reversible. The N–N distance in crystalline NO is 218 pm, nearly twice the N–O distance.[6]

Since the heat of formation of NO is endothermic, NO can be decomposed to the elements. Catalytic converters in cars exploit this reaction:

2 NO → O2 + N2

When exposed to oxygen, nitric oxide converts into nitrogen dioxide:

2 NO + O2 → 2 NO2

This reaction is thought to occur via the intermediates ONOO and the red compound ONOONO.[12]

In water, nitric oxide reacts with oxygen to form nitrous acid (HNO2). The reaction is thought to proceed via the following stoichiometry:

4 NO + O2 + 2 H2O → 4 HNO2

Nitric oxide reacts with fluorine, chlorine, and bromine to form the nitrosyl halides, such as nitrosyl chloride:

2 NO + Cl2 → 2 NOCl

With NO2, also a radical, NO combines to form the intensely blue dinitrogen trioxide:[6]

NO + NO2 ⇌ ON−NO2

Organic chemistry


The addition of a nitric oxide moiety to another molecule is often referred to as nitrosylation. The Traube reaction[13] is the addition of a two equivalents of nitric oxide onto an enolate, giving a diazeniumdiolate (also called a nitrosohydroxylamine).[14] The product can undergo a subsequent retro-aldol reaction, giving an overall process similar to the haloform reaction. For example, nitric oxide reacts with acetone and an alkoxide to form a diazeniumdiolate on each α position, with subsequent loss of methyl acetate as a by-product:[15]


This reaction, which was discovered around 1898, remains of interest in nitric oxide prodrug research. Nitric oxide can also react directly with sodium methoxide, ultimately forming sodium formate and nitrous oxide by way of an N-methoxydiazeniumdiolate.[16]

Coordination complexes


Nitric oxide reacts with transition metals to give complexes called metal nitrosyls. The most common bonding mode of nitric oxide is the terminal linear type (M−NO).[6] Alternatively, nitric oxide can serve as a one-electron pseudohalide. In such complexes, the M−N−O group is characterized by an angle between 120° and 140°. The NO group can also bridge between metal centers through the nitrogen atom in a variety of geometries.

Production and preparation


In commercial settings, nitric oxide is produced by the oxidation of ammonia at 750–900 °C (normally at 850 °C) with platinum as catalyst in the Ostwald process:

4 NH3 + 5 O2 → 4 NO + 6 H2O

The uncatalyzed endothermic reaction of oxygen (O2) and nitrogen (N2), which is effected at high temperature (>2000 °C) by lightning has not been developed into a practical commercial synthesis (see Birkeland–Eyde process):

N2 + O2 → 2 NO

Laboratory methods


In the laboratory, nitric oxide is conveniently generated by reduction of dilute nitric acid with copper:

8 HNO3 + 3 Cu → 3 Cu(NO3)2 + 4 H2O + 2 NO

An alternative route involves the reduction of nitrous acid in the form of sodium nitrite or potassium nitrite:

2 NaNO2 + 2 NaI + 2 H2SO4 → I2 + 2 Na2SO4 + 2 H2O + 2 NO
2 NaNO2 + 2 FeSO4 + 3 H2SO4 → Fe2(SO4)3 + 2 NaHSO4 + 2 H2O + 2 NO
3 KNO2 + KNO3 + Cr2O3 → 2 K2CrO4 + 4 NO

The iron(II) sulfate route is simple and has been used in undergraduate laboratory experiments. So-called NONOate compounds are also used for nitric oxide generation.

Detection and assay

Nitric oxide (white) in conifer cells, visualized using DAF-2 DA (diaminofluorescein diacetate)

Nitric oxide concentration can be determined using a chemiluminescent reaction involving ozone.[17] A sample containing nitric oxide is mixed with a large quantity of ozone. The nitric oxide reacts with the ozone to produce oxygen and nitrogen dioxide, accompanied with emission of light (chemiluminescence):

NO + O3NO2 + O2 +

which can be measured with a photodetector. The amount of light produced is proportional to the amount of nitric oxide in the sample.

Other methods of testing include electroanalysis (amperometric approach), where ·NO reacts with an electrode to induce a current or voltage change. The detection of NO radicals in biological tissues is particularly difficult due to the short lifetime and concentration of these radicals in tissues. One of the few practical methods is spin trapping of nitric oxide with iron-dithiocarbamate complexes and subsequent detection of the mono-nitrosyl-iron complex with electron paramagnetic resonance (EPR).[18][19]

A group of fluorescent dye indicators that are also available in acetylated form for intracellular measurements exist. The most common compound is 4,5-diaminofluorescein (DAF-2).[20]

Environmental effects


Acid rain deposition


Nitric oxide reacts with the hydroperoxyl radical (HO
) to form nitrogen dioxide (NO2), which then can react with a hydroxyl radical (HO) to produce nitric acid (HNO3):

NO2 + HO
NO2 + HO → HNO3

Nitric acid, along with sulfuric acid, contributes to acid rain deposition.

Ozone depletion


NO participates in ozone layer depletion. Nitric oxide reacts with stratospheric ozone to form O2 and nitrogen dioxide:

NO + O3NO2 + O2

This reaction is also utilized to measure concentrations of NO in control volumes.

Precursor to NO2


As seen in the acid deposition section, nitric oxide can transform into nitrogen dioxide (this can happen with the hydroperoxy radical, HO
, or diatomic oxygen, O2). Symptoms of short-term nitrogen dioxide exposure include nausea, dyspnea and headache. Long-term effects could include impaired immune and respiratory function.[21]

Biological functions


NO is a gaseous signaling molecule.[22] It is a key vertebrate biological messenger, playing a role in a variety of biological processes.[23] It is a bioproduct in almost all types of organisms, including bacteria, plants, fungi, and animal cells.[24]

Nitric oxide, an endothelium-derived relaxing factor (EDRF), is biosynthesized endogenously from L-arginine, oxygen, and NADPH by various nitric oxide synthase (NOS) enzymes.[25] Reduction of inorganic nitrate may also make nitric oxide.[26] One of the main enzymatic targets of nitric oxide is guanylyl cyclase.[27] The binding of nitric oxide to the heme region of the enzyme leads to activation, in the presence of iron.[27] Nitric oxide is highly reactive (having a lifetime of a few seconds), yet diffuses freely across membranes. These attributes make nitric oxide ideal for a transient paracrine (between adjacent cells) and autocrine (within a single cell) signaling molecule.[26] Once nitric oxide is converted to nitrates and nitrites by oxygen and water, cell signaling is deactivated.[27]

The endothelium (inner lining) of blood vessels uses nitric oxide to signal the surrounding smooth muscle to relax, resulting in vasodilation and increasing blood flow.[26] Sildenafil (Viagra) is a drug that uses the nitric oxide pathway. Sildenafil does not produce nitric oxide, but enhances the signals that are downstream of the nitric oxide pathway by protecting cyclic guanosine monophosphate (cGMP) from degradation by cGMP-specific phosphodiesterase type 5 (PDE5) in the corpus cavernosum, allowing for the signal to be enhanced, and thus vasodilation.[25] Another endogenous gaseous transmitter, hydrogen sulfide (H2S) works with NO to induce vasodilation and angiogenesis in a cooperative manner.[28][29]

Nasal breathing produces nitric oxide within the body, while oral breathing does not.[30][31]

Occupational safety and health


In the U.S., the Occupational Safety and Health Administration (OSHA) has set the legal limit (permissible exposure limit) for nitric oxide exposure in the workplace as 25 ppm (30 mg/m3) over an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of 25 ppm (30 mg/m3) over an 8-hour workday. At levels of 100 ppm, nitric oxide is immediately dangerous to life and health.[32]

Explosion hazard


Liquid nitrogen oxide is very sensitive to detonation even in the absence of fuel, and can be initiated as readily as nitroglycerin. Detonation of the endothermic liquid oxide close to its b.p. (-152°C) generated a 100 kbar pulse and fragmented the test equipment. It is the simplest molecule that is capable of detonation in all three phases. The liquid oxide is sensitive and may explode during distillation, and this has been the cause of industrial accidents.[33] Gaseous nitric oxide detonates at about 2300 m/s, but as a solid it can reach a detonation velocity of 6100 m/s.[34]




  1. ^ a b Nomenclature of Inorganic Chemistry, IUPAC Recommendations (PDF). International Union of Pure and Applied Chemistry. 2005. p. 69.
  2. ^ "Nitric Oxide (CHEBI:16480)". Chemical Entities of Biological Interest (ChEBI). UK: European Bioinformatics Institute.
  3. ^ a b c "Nitrogen monoxide - Registration Dossier - ECHA". Retrieved 2020-11-02.
  4. ^ a b c d "Safety Data Sheet - Nitric Oxide, compressed - Registration Dossier" (PDF). Retrieved 2020-11-02.
  5. ^ a b "Nitric oxide". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  6. ^ a b c d e Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  7. ^ Hou, Y. C.; Janczuk, A.; Wang, P. G. (1999). "Current trends in the development of nitric oxide donors". Current Pharmaceutical Design. 5 (6): 417–441. doi:10.2174/138161280506230110111042. PMID 10390607.
  8. ^ Culotta, Elizabeth; Koshland, Daniel E. Jr. (1992). "NO news is good news". Science. 258 (5090): 1862–1864. Bibcode:1992Sci...258.1862C. doi:10.1126/science.1361684. PMID 1361684.
  9. ^ "The Nobel Prize in Physiology or Medicine 1998". Retrieved 2022-06-17.
  10. ^ Berkowitz, Joseph (1979). Photoabsorption, Photoionization, and Photoelectron Spectroscopy. Academic Press. p. 231. doi:10.1016/B978-0-12-091650-4.50012-8.
  11. ^ Hoy, A. R.; Johns, J. W. C.; McKellar, A. R. W. (1975). "Stark Spectroscopy with the CO Laser: Dipole Moments, Hyperfine Structure, and Level Crossing Effects in the Fundamental Band of NO". Canadian Journal of Physics. 53 (19): 2029–2039. Bibcode:1975CaJPh..53.2029H. doi:10.1139/p75-254.
  12. ^ Galliker, Benedikt; et al. (2009). "Intermediates in the Autoxidation of Nitrogen Monoxide". Chemistry - A European Journal. 15 (25): 6161–6168. doi:10.1002/chem.200801819. ISSN 0947-6539. PMID 19437472.
  13. ^
  14. ^ Arulsamy, Navamoney; Bohle, D. Scott (2006). "Synthesis of Diazeniumdiolates from the Reactions of Nitric Oxide with Enolates". J. Org. Chem. 71 (2): 572–581. doi:10.1021/jo051998p. PMID 16408967.
  15. ^ Traube, Wilhelm (1898). "Ueber Synthesen stickstoffhaltiger Verbindungen mit Hülfe des Stickoxyds". Justus Liebig's Annalen der Chemie (in German). 300 (1): 81–128. doi:10.1002/jlac.18983000108.
  16. ^ Derosa, Frank; Keefer, Larry K.; Hrabie, Joseph A. (2008). "Nitric Oxide Reacts with Methoxide". The Journal of Organic Chemistry. 73 (3): 1139–1142. doi:10.1021/jo7020423. PMID 18184006.
  17. ^ Fontijn, Arthur; Sabadell, Alberto J.; Ronco, Richard J. (1970). "Homogeneous chemiluminescent measurement of nitric oxide with ozone. Implications for continuous selective monitoring of gaseous air pollutants". Analytical Chemistry. 42 (6): 575–579. doi:10.1021/ac60288a034.
  18. ^ Vanin, A; Huisman, A; Van Faassen, E (2002). "Iron dithiocarbamate as spin trap for nitric oxide detection: Pitfalls and successes". Nitric Oxide, Part D: Oxide Detection, Mitochondria and Cell Functions, and Peroxynitrite Reactions. Methods in Enzymology. Vol. 359. pp. 27–42. doi:10.1016/S0076-6879(02)59169-2. ISBN 9780121822620. PMID 12481557.
  19. ^ Nagano, T; Yoshimura, T (2002). "Bioimaging of nitric oxide". Chemical Reviews. 102 (4): 1235–1270. doi:10.1021/cr010152s. PMID 11942795.
  20. ^ Kojima H, Nakatsubo N, Kikuchi K, Kawahara S, Kirino Y, Nagoshi H, Hirata Y, Nagano T (1998). "Detection and imaging of nitric oxide with novel fluorescent indicators: diaminofluoresceins". Anal. Chem. 70 (13): 2446–2453. doi:10.1021/ac9801723. PMID 9666719.
  21. ^ "Centers for Disease Control and Prevention". NIOSH. 1 July 2014. Retrieved 10 December 2015.
  22. ^ Liu, Hongying; Weng, Lingyan; Yang, Chi (2017-03-28). "A review on nanomaterial-based electrochemical sensors for H2O2, H2S and NO inside cells or released by cells". Microchimica Acta. 184 (5): 1267–1283. doi:10.1007/s00604-017-2179-2. ISSN 0026-3672. S2CID 21308802.
  23. ^ Weller, Richard, Could the sun be good for your heart? Archived 2014-02-16 at the Wayback Machine TedxGlasgow. Filmed March 2012, posted January 2013
  24. ^ Roszer, T (2012) The Biology of Subcellular Nitric Oxide. ISBN 978-94-007-2818-9
  25. ^ a b Perez, Krystle M.; Laughon, Matthew (November 2015). "Sildenafil in Term and Premature Infants: A Systematic Review". Clinical Therapeutics. 37 (11): 2598–2607.e1. doi:10.1016/j.clinthera.2015.07.019. ISSN 0149-2918. PMID 26490498.
  26. ^ a b c Stryer, Lubert (1995). Biochemistry (4th ed.). W.H. Freeman and Company. p. 732. ISBN 978-0-7167-2009-6.
  27. ^ a b c T., Hancock, John (2010). Cell signalling (3rd ed.). Oxford: Oxford University Press. ISBN 9780199232109. OCLC 444336556.{{cite book}}: CS1 maint: multiple names: authors list (link)
  28. ^ Szabo, Csaba; Coletta, Ciro; Chao, Celia; Módis, Katalin; Szczesny, Bartosz; Papapetropoulos, Andreas; Hellmich, Mark R. (2013-07-23). "Tumor-derived hydrogen sulfide, produced by cystathionine-β-synthase, stimulates bioenergetics, cell proliferation, and angiogenesis in colon cancer". Proceedings of the National Academy of Sciences of the United States of America. 110 (30): 12474–12479. Bibcode:2013PNAS..11012474S. doi:10.1073/pnas.1306241110. ISSN 1091-6490. PMC 3725060. PMID 23836652.
  29. ^ Altaany, Zaid; Yang, Guangdong; Wang, Rui (July 2013). "Crosstalk between hydrogen sulfide and nitric oxide in endothelial cells". Journal of Cellular and Molecular Medicine. 17 (7): 879–888. doi:10.1111/jcmm.12077. ISSN 1582-4934. PMC 3822893. PMID 23742697.
  30. ^ Yasuda, Yoshifumi; Itoh, Tomonori; Miyamura, Miharu; Nishino, Hitoo (1997). "Comparison of Exhaled Nitric Oxide and Cardiorespiratory Indices between Nasal and Oral Breathing during Submaximal Exercise in Humans". Japanese Journal of Physiology. 47 (5): 465–470. doi:10.2170/jjphysiol.47.465. ISSN 0021-521X. PMID 9504133. Retrieved 2022-11-17.
  31. ^ Dahl, Melissa (2011-01-11). "'Mouth-breathing' gross, harmful to your health". NBC News. Retrieved 2021-09-06.
  32. ^ "Nitric oxide". National Institute for Occupational Safety and Health. Retrieved 2015-11-20.
  33. ^ Urben, Peter (22 May 2017). Bretherick's Handbook of Reactive Chemical Hazards | ScienceDirect. Elsevier Science. ISBN 9780081009710. Retrieved 2022-02-23.
  34. ^ Ribovich, John; Murphy, John; Watson, Richard (1975-01-01). "Detonation studies with nitric oxide, nitrous oxide, nitrogen tetroxide, carbon monoxide, and ethylene". Journal of Hazardous Materials. 1 (4): 275–287. Bibcode:1975JHzM....1..275R. doi:10.1016/0304-3894(75)80001-X. ISSN 0304-3894.

Further reading