Nitric oxide

From Wikipedia, the free encyclopedia
Jump to navigation Jump to search
Nitric oxide
Skeletal formula of nitric oxide with bond length
Skeletal formula showing three lone pairs and one unpaired electron
Space-filling model of nitric oxide
Names
IUPAC name
Nitrogen monoxide
Systematic IUPAC name
Oxidonitrogen(•)[1] (additive)
Other names
Nitric oxide
Nitrogen(II) oxide
Identifiers
3D model (JSmol)
3DMet B00122
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.030.233
EC Number 233-271-0
451
KEGG
RTECS number QX0525000
UNII
UN number 1660
Properties
NO
Molar mass 30.01 g·mol−1
Appearance Colourless gas
Density 1.3402 g dm−3
Melting point −164 °C (−263 °F; 109 K)
Boiling point −152 °C (−242 °F; 121 K)
0.0098 g/100ml (0 °C)
0.0056 g/100ml (20 °C)
1.0002697
Structure
linear (point group Cv)
Thermochemistry
210.76 J K−1 mol−1
91.29 kJ mol−1
Pharmacology
R07AX01 (WHO)
License data
Inhalation
Pharmacokinetics:
good
via pulmonary capillary bed
2–6 seconds
Hazards
Safety data sheet External MSDS
Oxidizing Agent O Toxic T
R-phrases (outdated) R8, R23, R34, R44
S-phrases (outdated) (S1), S17, S23, S36/37/39, S45
NFPA 704
Lethal dose or concentration (LD, LC):
315 ppm (rabbit, 15 min)
854 ppm (rat, 4 hr)
320 ppm (mouse)[2]
2500 ppm (mouse, 12 min)[2]
Related compounds
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 ☑Y☒N ?)
Infobox references

Nitric oxide (nitrogen oxide[3] or nitrogen monoxide) is a colorless gas with the formula NO. It is one of the principal oxides of nitrogen. Nitric oxide is a free radical, i.e., it has an unpaired electron, which is sometimes denoted by a dot in its chemical formula, i.e., ·NO. Nitric oxide is also a heteronuclear diatomic molecule, a historic class that drew researches which spawned early modern theories of chemical bonding.[4]

An important intermediate in chemical industry, 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.[5] It was proclaimed the "Molecule of the Year" in 1992.[6] The 1998 Nobel Prize in Physiology or Medicine was awarded for discovering nitric oxide's role as a cardiovascular signalling molecule.

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

Reactions[edit]

With di- and triatomic molecules[edit]

Upon condensing to a liquid, nitric oxide dimerizes, but the association is weak and reversible. The N---N distance in crystalline NO is 218 pm, nearly twice the N-O distance.[4]

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 conversion has been speculated as occurring via the ONOONO intermediate. In water, nitric oxide reacts with oxygen and water to form HNO2 or nitrous acid. 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:[4]

NO + NO2 ⇌ ON-NO2

Organic chemistry[edit]

Nitric oxide reacts with acetone and an alkoxide to a diazeniumdiolate or nitrosohydroxylamine and methyl acetate:[7]

Traube reaction

This reaction, which was discovered around 1898, remains of interest in nitric oxide prodrug research. Nitric oxide can also react directly with sodium methoxide, forming sodium formate and nitrous oxide.[8]

Coordination complexes[edit]

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).[4] 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[edit]

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

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[edit]

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 + 4 NaHSO4 + 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[edit]

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.[9] 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 + O3 → NO2 + 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).[10][11]

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).[12]

Environmental effects[edit]

Acid deposition[edit]

Nitric oxide reacts with the hydroperoxy radical (HO2) to form nitrogen dioxide (NO2), which then can react with a hydroxyl radical (OH) to produce nitric acid (HNO3):

·NO + HO2NO2 + OH
·NO2 + OH → HNO3

Nitric acid, along with sulfuric acid, contribute acid rain deposition.

Ozone depletion[edit]

Furthermore, ·NO participates in ozone layer depletion. In this process, nitric oxide reacts with stratospheric ozone to form O2 and nitrogen dioxide:

·NO + O3 → NO2 + O2

As seen in the Concentration Measurement section, this reaction is also utilized to measure concentrations of ·NO in control volumes.

Precursor to NO2[edit]

As seen in the Acid deposition section, nitric oxide can transform into nitrogen dioxide (this can happen with the hydroperoxy radical, HO2, 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.[13]

Biological functions[edit]

NO is a gaseous signaling molecule. It is a key vertebrate biological messenger, playing a role in a variety of biological processes.[14] It is a known bioproduct in almost all types of organisms, ranging from bacteria to plants, fungi, and animal cells.[15]

Nitric oxide, known as an endothelium-derived relaxing factor (EDRF), is biosynthesized endogenously from L-arginine, oxygen, and NADPH by various nitric oxide synthase (NOS) enzymes. Reduction of inorganic nitrate may also serve to make nitric oxide. The endothelium (inner lining) of blood vessels uses nitric oxide to signal the surrounding smooth muscle to relax, thus resulting in vasodilation and increasing blood flow. 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.[16]

Occupational safety and health[edit]

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.[17]

References[edit]

  1. ^ "Nitric Oxide (CHEBI:16480)". Chemical Entities of Biological Interest (ChEBI). UK: European Bioinformatics Institute. 
  2. ^ a b "Nitric oxide". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH). 
  3. ^ IUPAC nomenclature of inorganic chemistry 2005. PDF.
  4. ^ a b c d Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  5. ^ Hou, YC; Janczuk, A; Wang, PG (1999). "Current trends in the development of nitric oxide donors". Current Pharmaceutical Design. 5 (6): 417–41. PMID 10390607. 
  6. ^ Culotta, Elizabeth; Koshland, Daniel E. Jr (1992). "NO news is good news". Science. 258 (5090): 1862–1864. doi:10.1126/science.1361684. PMID 1361684. 
  7. ^ Traube, Wilhelm (1898). "Ueber Synthesen stickstoffhaltiger Verbindungen mit Hülfe des Stickoxyds". Justus Liebig's Annalen der Chemie. 300: 81–128. doi:10.1002/jlac.18983000108. 
  8. ^ Derosa, Frank; Keefer, Larry K.; Hrabie, Joseph A. (2008). "Nitric Oxide Reacts with Methoxide". The Journal of Organic Chemistry. 73 (3): 1139–42. doi:10.1021/jo7020423. PMID 18184006. 
  9. ^ 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. 
  10. ^ Vanin, A; Huisman, A; Van Faassen, E (2002). "Iron dithiocarbamate as spin trap for nitric oxide detection: Pitfalls and successes". Methods in Enzymology. Methods in Enzymology. 359: 27–42. doi:10.1016/S0076-6879(02)59169-2. ISBN 9780121822620. PMID 12481557. 
  11. ^ Nagano, T; Yoshimura, T (2002). "Bioimaging of nitric oxide". Chemical Reviews. 102 (4): 1235–70. doi:10.1021/cr010152s. PMID 11942795. 
  12. ^ 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. 
  13. ^ "Centers for Disease Control and Prevention". NIOSH. 1 July 2014. Retrieved 10 December 2015. 
  14. ^ Weller, Richard, Could the sun be good for your heart? TedxGlasgow. Filmed March 2012, posted January 2013
  15. ^ Roszer, T (2012) The Biology of Subcellular Nitric Oxide. ISBN 978-94-007-2818-9
  16. ^ Stryer, Lubert (1995). Biochemistry, 4th Edition. W.H. Freeman and Company. p. 732. ISBN 0-7167-2009-4. 
  17. ^ "CDC - NIOSH Pocket Guide to Chemical Hazards - Nitric oxide". www.cdc.gov. Retrieved 2015-11-20. 

Further reading[edit]

External links[edit]