 | |||
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| Names | |||
|---|---|---|---|
| IUPAC name Nitrogen monoxide[1] | |||
| Systematic IUPAC name Oxidonitrogen(•)[2] (additive) | |||
| Other names Nitrogen oxide Nitrogen(II) oxide Oxonitrogen Nitrogen monoxide | |||
| Identifiers | |||
3D model (JSmol) | |||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider |
| ||
| DrugBank |
| ||
| ECHA InfoCard | 100.030.233 | ||
| EC Number |
| ||
| 451 | |||
| KEGG |
| ||
| RTECS number |
| ||
| UNII | |||
| UN number | 1660 | ||
| |||
| |||
| Properties | |||
| NO | |||
| 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) | |||
Refractive index (nD) | 1.0002697 | ||
| Structure | |||
| linear (point group C∞v) | |||
| Thermochemistry | |||
Std molar entropy(S⦵298) | 210.76 J/(K·mol) | ||
Std enthalpy of formation(ΔfH⦵298) | 90.29 kJ/mol | ||
| Pharmacology | |||
| R07AX01 (WHO) | |||
| License data | |||
| Inhalation | |||
| Pharmacokinetics: | |||
| good | |||
| via pulmonary capillary bed | |||
| 2–6 seconds | |||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards | Very toxic, corrosive, oxidizer[4] | ||
| GHS labelling: | |||
   [3][4] | |||
| Danger | |||
| H270,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): | |||
LC50 (median concentration) | 315 ppm (rabbit, 15 min) 854 ppm (rat, 4 h) 2500 ppm (mouse, 12 min)[5] | ||
LCLo (lowest published) | 320 ppm (mouse)[5] | ||
| Safety data sheet (SDS) | External SDS | ||
| Related compounds | |||
Related nitrogen oxides | Dinitrogen pentoxide Dinitrogen tetroxide | ||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Nitric oxide (nitrogen oxide ornitrogen monoxide[1]) is a colorless gas with the formulaNO. It is one of the principaloxides of nitrogen. Nitric oxide is afree radical: it has anunpaired electron, which is sometimes denoted by a dot in itschemical formula (•N=O or•NO). Nitric oxide is also aheteronucleardiatomic molecule, a class of molecules whose study spawned early moderntheories of chemical bonding.[6]
An importantintermediate inindustrial chemistry, nitric oxide forms in combustion systems and can be generated by lightning in thunderstorms. In mammals, including humans, nitric oxide is asignaling molecule in many physiological and pathological processes.[7] It was proclaimed the "Molecule of the Year" in 1992.[8] The1998 Nobel Prize in Physiology or Medicine was awarded for discovering nitric oxide's role as a cardiovascular signalling molecule.[9] Its impact extends beyond biology, with applications in medicine, such as the development ofsildenafil (Viagra), and in industry, includingsemiconductor manufacturing.[10][11]
Nitric oxide should not be confused withnitrogen dioxide (NO2), a brown gas and majorair pollutant, or withnitrous oxide (N2O), ananesthetic gas.[6]
Nitric oxide (NO) was first identified byJoseph Priestley in the late 18th century, originally seen as merely a toxic byproduct of combustion and an environmental pollutant.[12] Its biological significance was later uncovered in the 1980s when researchersRobert F. Furchgott,Louis J. Ignarro, andFerid Murad discovered its critical role as avasodilator in the cardiovascular system, a breakthrough that earned them the 1998 Nobel Prize in Physiology or Medicine.[13]
The ground state electronic configuration of NO is, in united atom notation:[14]
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 momentumJ=3⁄2 orJ=1⁄2.
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.[15]
Upon condensing to aneat liquid, nitric oxidedimerizes to colorlessdinitrogen dioxide (O=N–N=O), but the association is weak and reversible. The N–N distance in crystalline NO is 218 pm, nearly twice the N–O distance. Condensation in a highly polar environment instead gives the red alternant isomer O=N–O=N.[6]
Since the heat of formation of•NO isendothermic, NO can be decomposed to the elements.Catalytic converters in cars exploit this reaction:
When exposed tooxygen, nitric oxide converts intonitrogen dioxide:
This reaction is thought to occur via the intermediates ONOO• and the red compound ONOONO.[16]
In water, nitric oxide reacts with oxygen to formnitrous acid (HNO2). The reaction is thought to proceed via the followingstoichiometry:
Nitric oxide reacts withfluorine,chlorine, andbromine to form the nitrosyl halides, such asnitrosyl chloride:
With NO2, also a radical, NO combines to form the intensely bluedinitrogen trioxide:[6]
Nitric oxide rarely sees organic chemistry use. Most reactions with it produce complex mixtures of salts, separable only through carefulrecrystallization.[17]
The addition of a nitric oxidemoiety to another molecule is often referred to asnitrosylation. TheTraube reaction is theaddition of a twoequivalents of nitric oxide onto anenolate, giving a diazeniumdiolate (also called anitrosohydroxylamine).[18] The product can undergo a subsequent retro-aldol reaction, giving an overall process similar to thehaloform reaction. For example, nitric oxide reacts withacetone and analkoxide to form a diazeniumdiolate on eachα position, with subsequent loss ofmethyl acetate as aby-product:[19]
This reaction, which was discovered around 1898, remains of interest in nitric oxideprodrug research. Nitric oxide can also react directly withsodium methoxide, ultimately formingsodium formate andnitrous oxide by way of anN-methoxydiazeniumdiolate.[20]
Sufficiently basicsecondary amines undergo a Traube-like reaction to giveNONOates.[21] However, very few nucleophiles undergo the Traube reaction, either failing to adduce NO or immediately decomposing withnitrous oxide release.[17]
Nitric oxide reacts withtransition metals to give complexes calledmetal 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.
In commercial settings, nitric oxide is produced by theoxidation ofammonia at 750–900 °C (normally at 850 °C) withplatinum ascatalyst in theOstwald process:
The uncatalyzedendothermic reaction ofoxygen (O2) andnitrogen (N2), which is effected at high temperature (>2000 °C) by lightning has not been developed into a practical commercial synthesis (seeBirkeland–Eyde process):
In the laboratory, nitric oxide is conveniently generated by reduction of dilutenitric acid withcopper:
An alternative route involves the reduction of nitrous acid in the form ofsodium nitrite orpotassium nitrite:
The iron(II) sulfate route is simple and has been used in undergraduate laboratory experiments.
So-calledNONOate compounds are also used for nitric oxide generation, especially in biological laboratories. However, other Traube adducts may decompose to instead givenitrous oxide.[22]
_(white)_in_the_cytoplasm_(green)_of_clusters_of_conifer_cells_one_hour_after_mechanical_agitation.jpg)
Nitric oxide concentration can be determined using achemiluminescent reaction involvingozone.[23] A sample containing nitric oxide is mixed with a large quantity of ozone. The nitric oxide reacts with the ozone to produceoxygen andnitrogen dioxide, accompanied with emission oflight (chemiluminescence):
which can be measured with aphotodetector. The amount of light produced is proportional to the amount of nitric oxide in the sample.
Other methods of testing includeelectroanalysis (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 isspin trapping of nitric oxide with iron-dithiocarbamate complexes and subsequent detection of the mono-nitrosyl-iron complex withelectron paramagnetic resonance (EPR).[24][25]
A group offluorescent dye indicators that are also available inacetylated form for intracellular measurements exist. The most common compound is4,5-diaminofluorescein (DAF-2).[26]
Nitric oxide reacts with thehydroperoxyl radical (HO•
2) to form nitrogen dioxide (NO2), which then can react with ahydroxyl radical (HO•) to producenitric acid (HNO3):
Nitric acid, along withsulfuric acid, contributes toacid rain deposition.
•NO participates inozone layer depletion. Nitric oxide reacts with stratosphericozone to form O2 and nitrogen dioxide:
This reaction is also utilized to measure concentrations of•NO in control volumes.
As seen in theacid deposition section, nitric oxide can transform into nitrogen dioxide (this can happen with the hydroperoxy radical,HO•
2, or diatomic oxygen, O2). Symptoms of short-term nitrogen dioxide exposure include nausea,dyspnea and headache. Long-term effects could include impaired immune andrespiratory function.[27]
NO is agaseous signaling molecule.[28] It is a keyvertebratebiological messenger, playing a role in a variety of biological processes.[29] It is a bioproduct in almost all types of organisms, including bacteria, plants, fungi, and animal cells.[30]
Nitric oxide, anendothelium-derived relaxing factor (EDRF), is biosynthesized endogenously fromL-arginine,oxygen, andNADPH by variousnitric oxide synthase (NOS)enzymes.[31] Reduction of inorganic nitrate may also make nitric oxide.[32] One of the main enzymatic targets of nitric oxide isguanylyl cyclase.[33] The binding of nitric oxide to theheme region of the enzyme leads to activation, in the presence of iron.[33] 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 transientparacrine (between adjacent cells) andautocrine (within a single cell) signaling molecule.[32] Once nitric oxide is converted to nitrates and nitrites by oxygen and water, cell signaling is deactivated.[33]
Theendothelium (inner lining) ofblood vessels uses nitric oxide to signal the surroundingsmooth muscle to relax, resulting invasodilation and increasing blood flow.[32]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 protectingcyclic guanosine monophosphate (cGMP) from degradation bycGMP-specific phosphodiesterase type 5 (PDE5) in thecorpus cavernosum, allowing for the signal to be enhanced, and thusvasodilation.[31] Another endogenous gaseous transmitter,hydrogen sulfide (H2S) works with NO to induce vasodilation and angiogenesis in a cooperative manner.[34][35]
Nasal breathing produces nitric oxide within the body, whileoral breathing does not.[36][37]
In the U.S., theOccupational 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. TheNational Institute for Occupational Safety and Health (NIOSH) has set arecommended exposure limit (REL) of 25 ppm (30 mg/m3) over an 8-hour workday. At levels of 100 ppm, nitric oxide isimmediately dangerous to life and health.[38]
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.[39] Gaseous nitric oxide detonates at about 2300 m/s, but as a solid it can reach a detonation velocity of 6100 m/s.[40]
Notes
