Nitric oxide

Nitric oxide
Nitric oxide Nitric oxide
Molecular formula NO
Molar mass 30.0061 g/mol
Appearance colourless gas
CAS number 10102-43-9
Density and phase 1.3 × 103 kg m−3 (liquid)
1.34 g dm−3 (vapour)
Solubility in water
Melting point −163.6 °C (109.6 K)
Boiling point −151.7 °C (121.4 K)
Molecular shape linear
Dipole moment 0.15 D
Thermodynamic data
Std enthalpy of
+90.2 kJ/mol
MSDS External MSDS
EU classification Toxic (T), corrosive (C)
NFPA 704

NFPA 704.svg

R-phrases R23, R24, R25, R34, R44
S-phrases S23, S36, S37, S39
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Related compounds
Related nitrogen oxides Nitrous oxide
Nitrogen dioxide
Dinitrogen trioxide
Dinitrogen tetroxide
Dinitrogen pentoxide
Related compounds Nitric acid
Nitrous acid
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)

The chemical compound nitric oxide is a gas with formula NO. It is an important signaling molecule in the body of mammals, including humans—one of the few gaseous signaling molecules known. It is also a toxic air pollutant produced by automobile engines and power plants.

Nitric oxide (NO) should not be confused with nitrous oxide (N2O), a general anesthetic, or with nitrogen dioxide (NO2), which is another poisonous air pollutant.


The nitric oxide molecule is a free radical, which is relevant to understanding its high reactivity. It reacts with the oxygen in air to form nitrogen dioxide, signaled by the appearance of the reddish-brown color.

Production and environmental effects

From a thermodynamic perspective, NO is unstable with respect to O2 and N2, although this conversion is very slow at ambient temperatures in the absence of a catalyst. Because the heat of formation of NO is endothermic, its synthesis from molecular nitrogen and oxygen requires elevated temperatures, >1,000 °C. A major natural source is lightning. The use of internal combustion engines has drastically increased the presence of nitric oxide in the environment. One purpose of catalytic converters in cars is to minimize NO formation by catalytic reversion to O2 and N2.

Nitric oxide in the air may convert to nitric acid, which has been implicated in acid rain. Furthermore, both NO and NO2 participate in ozone layer depletion. Nitric oxide (NO) is a small highly diffusible gas and a ubiquitous bioactive molecule.

Technical applications

Although NO has relatively few direct uses, it is produced on a massive scale as an intermediate in the Ostwald process for the synthesis of nitric acid from ammonia. In 2005, the United States alone produced six million metric tons of nitric acid.[1] It finds use in the semiconductor industry for various processes. In one of its applications it is used along with nitrous oxide to form oxynitride gates in CMOS devices.

Nitric oxide can also be used for detecting surface radicals on polymers. Quenching of surface radicals with nitric oxide results in the incorporation of nitrogen, which can be quantified by means of X-ray photoelectron spectroscopy.

Biological functions

Nitric oxide is a key biological messenger, playing a role in a variety of biological process. Nitric oxide, known as the endothelium-derived relaxing factor, or EDRF, is biosynthesized from arginine and oxygen by various nitric oxide synthase (NOS) enzymes and by reduction of inorganic nitrate. The endothelium (inner lining) of blood vessels use nitric oxide to signal the surrounding smooth muscle to relax, thus dilating the artery and increasing blood flow. The production of nitric oxide is elevated in populations living at high-altitudes, which helps these people avoid hypoxia. Effects include blood vessel dilatation, neurotransmission, modulation of the hair cycle, and penile erections. Nitroglycerin and amyl nitrite serve as vasodilators because they are converted to nitric oxide in the body.

Nitric oxide is also generated by macrophages and neutrophils as part of the human immune response. Nitric oxide is toxic to bacteria and other human pathogens. Many bacterial pathogens have evolved mechanisms for nitric oxide resistance.[2]

Nitric oxide can contribute to reperfusion injury when excessive amount produced during reperfusion (following a period of ischemia) reacts with superoxide to produce the damaging free radical peroxynitrite. Inhaled nitric oxide has been shown to help survival and recovery from paraquat poisoning, which produces lung tissue damaging superoxide and hinders NOS metabolism.

In plants, nitric oxide can be produced by any of four routes: (i) nitric oxide synthase (as in animals), (ii) by plasma membrane-bound nitrate reductase, (iii) by mitochondrial electron transport chain, or (iv) by non enzymatic reactions. It is a signaling molecule, acts mainly against oxidative stress and also plays a role in plant pathogen interactions. Treating cut flowers and other plants with nitric oxide has been shown to lengthen the time before wilting.[3]

A biologically important reaction of nitric oxide is S-nitrosation (or S-nitrosylation), the covalent attachment of a nitric oxide to the thiol group of cysteine within proteins. S-Nitrosylation has been described by some of its proponents as a mechanism for dynamic, post-translational regulation of most or all main classes of protein. Firm evidence to support this claim is limited.


When exposed to oxygen, NO is converted into NO2.

2NO + O2 → 2NO2

This conversion has been speculated as occurring via the ONOONO intermediate. In water, NO react 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

NO will react with fluorine, chlorine, and bromine to from the XNO species, known as the nitrosyl halides, such as nitrosyl chloride. Nitrosyl iodide can form but is an extremely short lived species and tends to reform I2.

2NO + Cl2 → 2NOCl

Nitroxyl (HNO) is the reduced form of nitric oxide.


As noted above, nitric oxide is produced industrially by the direct reaction of O2 and N2 at high temperatures. In the laboratory, it is conveniently generated by the reduction of nitric acid:

8HNO3 + 3Cu → 3Cu(NO3)2 + 4H2O + 2NO

It can also be prepared by the reduction of nitrous acid:

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(l) + KNO3 (l) + Cr2O3(s) → 2 K2CrO4(s) + 4 NO (g)

The route involving iron(II) sulfate is simple and has been used in undergraduate laboratory experiments.

Commercially, NO is produced by the oxidation of ammonia at 750 to 900 °C (normally at 850 °C) in the presence of platinum as catalyst:

4NH3 + 5O2 → 4NO + 6H2O

The uncatalyzed endothermic reaction of O2 and N2, which occurs at high temperatures (above 2,000 °C) with lightning, has not been developed into a practical commercial synthesis. The reaction may be written as:

N2 + O2 → 2NO

Coordination Chemistry

Nitric oxide forms complexes with all transition metals to give complexes called metal nitrosyls. The most common bonding mode of NO is the terminal "linear" type, in which the angle of the M-N-O group varies from 160 to 180 degrees. In this case, the NO group is formally considered a 3-electron donor. Alternatively, one can view such complexes as derived from NO+, which is isoelectronic with CO.

In other complexes, the M-N-O group is characterized by an angle between 120-140 degrees. In such cases, nitric oxide is thought of as a "one-electron pseudohalide."

The NO group can also form a bridge between metal centers through the nitrogen atom in a variety of geometries.

Measurement of nitric oxide concentration

The concentration of nitric oxide can be determined using a simple chemiluminescent reaction involving ozone: 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. This reaction also produces light (chemiluminescence), which can be measured with a photodetector. The amount of light produced is proportional to the amount of nitric oxide in the sample.

NO + O3 → NO2 + O2 + light

Other methods of testing include electroanalysis, where NO reacts with an electrode to induce a current or voltage change.


  1. "Production: Growth is the Norm." Chemical and Engineering News (July 10, 2006): 59.
  2. Janeway. C. A., et al. Immunobiology: the immune system in health and disease 6th ed. New York: Garland Science, 2005. ISBN 0815341016
  3. Siegel-Itzkovich, Judy. "Viagra makes flowers stand up straight." Student BMJ (September 7, 1999). Retrieved May 20, 2007.


  • Butler, A. R., and R. Nicholson. 2003. Life, Death and Nitric Oxide. Cambridge, UK: Royal Society of Chemistry. ISBN 0854046860
  • Cotton, F. Albert, Geoffrey Wilkinson, Carlos A. Murillo, and Manfred Bochmann. 1999. Advanced Inorganic Chemistry. 6th edition. New York: Wiley. ISBN 0471199575
  • Faassen, Ernst van and Anatoly F. Vanin (eds.). 2007. Radicals for Life: The Various Forms of Nitric Oxide. Amsterdam: Elsevier. ISBN 0444522360
  • Lancaster, Jack, Jr. 1996. Nitric Oxide: Principles and Actions. San Diego, CA: Academic Press. ISBN 0124355552

External links

All links retrieved January 15, 2015.


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