Nitrogen dioxide

Nitrogen dioxide is the chemical compound with the formula NO
2
. It is one of several nitrogen oxides. NO
2
is an intermediate in the industrial synthesis of nitric acid, millions of tons of which are produced each year which is used primarily in the production of fertilizers. At higher temperatures it is a reddish-brown gas that has a characteristic sharp, biting odor and is a prominent air pollutant.[8] Nitrogen dioxide is a paramagnetic, bent molecule with C2v point group symmetry.

Nitrogen dioxide
Skeletal formula of nitrogen dioxide with some measurementsEP
Spacefill model of nitrogen dioxide
Nitrogen dioxide at different temperatures

Nitrogen dioxide at −196 °C, 0 °C, 23 °C, 35 °C, and 50 °C. (NO
2
) converts to the colorless dinitrogen tetroxide (N
2
O
4
) at low temperatures, and reverts to NO
2
at higher temperatures.
Names
IUPAC name
Nitrogen dioxide
Other names
Nitrogen(IV) oxide,[1] Deutoxide of nitrogen
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.234
EC Number 233-272-6
976
RTECS number QW9800000
UNII
UN number 1067
Properties
NO
2
Molar mass 46.006 g mol−1[2]
Appearance Brown gas[2]
Odor Chlorine like
Density 1.880 g dm−3[2]
Melting point −9.3 °C (15.3 °F; 263.8 K)[2]
Boiling point 21.15 °C (70.07 °F; 294.30 K)[2]
Hydrolyses
Solubility soluble in CCl
4
, nitric acid,[3] chloroform
Vapor pressure 98.80 kPa (at 20 °C)
+150.0·10−6 cm3/mol[4]
1.449 (at 20 °C)
Structure
C2v
Bent
Thermochemistry[5]
37.2 J/mol K
240.1 J mol−1 K−1
+33.2 kJ mol−1
Hazards
Main hazards Poison, oxidizer
Safety data sheet ICSC 0930
GHS pictograms The flame-over-circle pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The gas-cylinder pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The skull-and-crossbones pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The health hazard pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
GHS signal word Danger
H270, H314, H330
P220, P260, P280, P284, P305+351+338, P310
NFPA 704
Lethal dose or concentration (LD, LC):
30 ppm (guinea pig, 1 hr)
315 ppm (rabbit, 15 min)
68 ppm (rat, 4 hr)
138 ppm (rat, 30 min)
1000 ppm (mouse, 10 min)[7]
64 ppm (dog, 8 hr)
64 ppm (monkey, 8 hr)[7]
US health exposure limits (NIOSH):
PEL (Permissible)
C 5 ppm (9 mg/m3)[6]
REL (Recommended)
ST 1 ppm (1.8 mg/m3)[6]
IDLH (Immediate danger)
20 ppm[6]
Related compounds
Dinitrogen pentoxide

Dinitrogen tetroxide
Dinitrogen trioxide
Nitric oxide
Nitrous oxide

Related compounds
Chlorine dioxide
Carbon dioxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Properties

Nitrogen dioxide is a reddish-brown gas above 21.2 °C (70.2 °F; 294.3 K) with a pungent, acrid odor, becomes a yellowish-brown liquid below 21.2 °C (70.2 °F; 294.3 K), and converts to the colorless dinitrogen tetroxide (N
2
O
4
) below −11.2 °C (11.8 °F; 261.9 K).[6]

The bond length between the nitrogen atom and the oxygen atom is 119.7 pm. This bond length is consistent with a bond order between one and two.

Unlike ozone, O3, the ground electronic state of nitrogen dioxide is a doublet state, since nitrogen has one unpaired electron,[9] which decreases the alpha effect compared with nitrite and creates a weak bonding interaction with the oxygen lone pairs. The lone electron in NO
2
also means that this compound is a free radical, so the formula for nitrogen dioxide is often written as NO
2
.

The reddish-brown color is a consequence of preferential absorption of light in the blue (400 – 500 nm), although the absorption extends throughout the visible (at shorter wavelengths) and into the infrared (at longer wavelengths). Absorption of light at wavelengths shorter than about 400 nm results in photolysis (to form NO + O, atomic oxygen); in the atmosphere the addition of O atom so formed to O2 results in ozone formation.

Preparation and reactions

Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:[10]

2 NO + O
2
→ 2 NO
2

Nitrogen dioxide is formed in most combustion processes using air as the oxidant. At elevated temperatures nitrogen combines with oxygen to form nitric oxide:

O
2
+ N
2
→ 2 NO

In the laboratory, NO
2
can be prepared in a two-step procedure where dehydration of nitric acid produces dinitrogen pentoxide, which subsequently undergoes thermal decomposition:

HNO
3
N
2
O
5
+ H
2
O
N
2
O
5
→ 4 NO
2
+ O
2

The thermal decomposition of some metal nitrates also affords NO
2
:

Pb(NO
3
)
2
→ 2 PbO + 4 NO
2
+ O
2

Alternatively, reduction of concentrated nitric acid by metal (such as copper).

HNO
3
+ Cu → Cu(NO
3
)
2
+ 2 NO
2
+ 2 H
2
O

Or finally by adding concentrated nitric acid over tin; hydrated tin dioxide is produced as byproduct.

4 HNO3 + Sn → H2O + H2SnO3 + 4 NO2

Main reactions

Basic thermal properties

NO
2
exists in equilibrium with the colourless gas dinitrogen tetroxide (N
2
O
4
):

NO
2
N
2
O
4

The equilibrium is characterized by ΔH = −57.23 kJ/mol, which is exothermic. NO2 is favored at higher temperatures, while at lower temperatures, dinitrogen tetroxide (N2O4) predominates. Dinitrogen tetroxide (N
2
O
4
) can be obtained as a white solid with melting point −11.2 °C.[10] NO2 is paramagnetic due to its unpaired electron, while N2O4 is diamagnetic.

The chemistry of nitrogen dioxide has been investigated extensively. At 150 °C, NO
2
decomposes with release of oxygen via an endothermic process (ΔH = 14 kJ/mol):

NO
2
→ 2 NO + O
2

As an oxidizer

As suggested by the weakness of the N–O bond, NO
2
is a good oxidizer. Consequently, it will combust, sometimes explosively, with many compounds, such as hydrocarbons.

Hydrolysis

It hydrolyses to give nitric acid and nitrous acid:

NO
2
(N
2
O
4
) + H
2
O
HNO
2
+ HNO
3

This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia.[11] This reaction is negligibly slow at low concentrations of NO2 characteristic of the ambient atmosphere, although it does proceed upon NO2 uptake to surfaces. Such surface reaction is thought to produce gaseous HNO2 (often written as HONO) in outdoor and indoor environments.[12]

Formation from decomposition of nitric acid

Nitric acid decomposes slowly to nitrogen dioxide by the overall reaction:

HNO
3
→ 4 NO
2
+ 2 H
2
O
+ O
2

The nitrogen dioxide so formed confers the characteristic yellow color often exhibited by this acid.

Conversion to nitrates

NO
2
is used to generate anhydrous metal nitrates from the oxides:[10]

MO + 3 NO
2
M(NO
3
)
2
+ NO

Conversion to nitrites

Alkyl and metal iodides give the corresponding nitrites:

CH
3
I
+ 2 NO
2
→ 2 CH
3
NO
2
+ I
2
TiI
4
+ 4 NO
2
Ti(NO
2
)
4
+ 2 I
2

Ecology

NO
2
is introduced into the environment by natural causes, including entry from the stratosphere, bacterial respiration, volcanos, and lightning. These sources make NO
2
a trace gas in the atmosphere of Earth, where it plays a role in absorbing sunlight and regulating the chemistry of the troposphere, especially in determining ozone concentrations.[13]

Uses

NO
2
is used as an intermediate in the manufacturing of nitric acid, as a nitrating agent in manufacturing of chemical explosives, as a polymerization inhibitor for acrylates, as a flour bleaching agent.,[14]:223 and as a room temperature sterilization agent.[15] It is also used as an oxidizer in rocket fuel, for example in red fuming nitric acid; it was used in the Titan rockets, to launch Project Gemini, in the maneuvering thrusters of the Space Shuttle, and in unmanned space probes sent to various planets.[16]

Human-caused sources and exposure

For the general public, the most prominent sources of NO
2
are internal combustion engines burning fossil fuels.[8] Outdoors, NO
2
can be a result of traffic from motor vehicles.[17]

Indoors, exposure arises from cigarette smoke,[18] and butane and kerosene heaters and stoves.[19]

Workers in industries where NO
2
is used are also exposed and are at risk for occupational lung diseases, and NIOSH has set exposure limits and safety standards.[6] Astronauts in the Apollo–Soyuz Test Project were almost killed when NO
2
was accidentally vented into the cabin.[16] Agricultural workers can be exposed to NO
2
arising from grain decomposing in silos; chronic exposure can lead to lung damage in a condition called "Silo-filler's disease".[20][21]

Historically, nitrogen dioxide was also produced by atmospheric nuclear tests, and was responsible for the reddish colour of mushroom clouds.[22]

Toxicity

Gaseous NO
2
diffuses into the epithelial lining fluid (ELF) of the respiratory epithelium and dissolves, and chemically reacts with antioxidant and lipid molecules in the ELF; NO
2
's health effects are caused by the reaction products or their metabolites, which are reactive nitrogen species and reactive oxygen species that can drive bronchoconstriction, inflammation, reduced immune response, and may have effects on the heart.[23]

No2toxpathwaysEPA
Pathways indicated by a dotted line are those for which evidence is limited to findings from experimental animal studies, while evidence from controlled human exposure studies is available for pathways indicated by a solid line. Dashed lines indicate proposed links to the outcomes of asthma exacerbation and respiratory tract infections. Key events are subclinical effects, endpoints are effects that are generally measured in the clinic, and outcomes are health effects at the organism level. NO2 = nitrogen dioxide; ELF = epithelial lining fluid.[23]:4–62
AirQualityLondon1
Nitrogen dioxide diffusion tube for air quality monitoring. Positioned in the City of London

Acute harm due to NO
2
exposure is only likely to arise in occupational settings. Direct exposure to the skin can cause irritations and burns. Only very high concentrations of the gaseous form cause immediate distress: 10–20 ppm can cause mild irritation of the nose and throat, 25–50 ppm can cause edema leading to bronchitis or pneumonia, and levels above 100 ppm can cause death due to asphyxiation from fluid in the lungs. There are often no symptoms at the time of exposure other than transient cough, fatigue or nausea, but over hours inflammation in the lungs causes edema.[24][25]

For skin or eye exposure, the affected area is flushed with saline. For inhalation, oxygen is administered, bronchodilators may be administered, and if there are signs of methemoglobinemia, a condition that arises when nitrogen-based compounds affect the hemoglobin in red blood cells, methylene blue may be administered.[26][27]

It is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and it is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.[28]

Health effects of NO
2
exposure

For the public, chronic exposure to NO
2
can cause respiratory effects including airway inflammation in healthy people and increased respiratory symptoms in people with asthma. NO
2
creates ozone which causes eye irritation and exacerbates respiratory conditions, leading to increased visits to emergency departments and hospital admissions for respiratory issues, especially asthma.[29]

The effects of toxicity on health have been examined using questionnaires and inperson interviews in an effort to understand the relationship between (NO
2
) and asthma. The influence of indoor air pollutants on health is important because the majority of people in the world spend more than 80% of their time indoors.[30] The amount of time spent indoors depends upon on several factors including geographical region, job activities, and gender among other variables. Additionally, because home insulation is improving, this can result in greater retention of indoor air pollutants, such as (NO
2
) .[30] With respect to geographic region, the prevalence of asthma has ranged from 2 to 20% with no clear indication as to what's driving the difference.[30] This may be a result of the “hygiene hypothesis” or "western lifestyle” that captures the notions of homes that are well insulated and with fewer inhabitants.[30] Another study examined the relationship between nitrogen exposure in the home and respiratory symptoms and found a statistically significant odds ratio of 2.23 (95% CI: 1.06, 4.72) among those with a medical diagnosis of asthma and gas stove exposure.[31]

A major source of indoor exposure to (NO
2
) is from the use of gas stoves for cooking or heating in homes. According to the 2000 census, over half of US households use gas stoves[32] and indoor exposure levels of (NO
2
) are, on average, at least three times higher in homes with gas stoves compared to electric stoves with the highest levels being in multifamily homes. Exposure to (NO
2
) is especially harmful for children with asthma. Research has shown that children with asthma who live in homes with gas stoves have greater risk of respiratory symptoms such as wheezing, cough and chest tightness.[31][33] Additionally, gas stove use was associated with reduced lung function in girls with asthma, although this association was not found in boys.[34] Using ventilation when operating gas stoves may reduce the risk of respiratory symptoms in children with asthma.

In a cohort study with inner-city minority African American Baltimore children to determine if there was a relationship between (NO
2
) and asthma for children aged 2 to 6 years old, with an existing medical diagnosis of asthma, and one asthma related visit. Families of lower socioeconomic status were more likely to have gas stoves in their homes. The study concluded that higher levels of (NO
2
) within a home were linked to a greater level of respiratory symptoms among the study population. This further exemplifies that (NO
2
) toxicity is dangerous for children.[35]

Avoiding NO
2
toxicity

While using a gas stove, it is advised to also use ventilation. Studies show that in homes with gas stoves, if ventilation is used while using gas stoves, then children have lower odds of asthma, wheezing and bronchitis as compared to children in homes that never used ventilation.[36] If venting isn't possible, then replacing gas stoves with electric stove could be another option. Replacing gas stoves with electric ranges could greatly reduce the exposure to indoor NO2 and improve the respiratory function of children with asthma. It is important to keep gas stoves and heaters in good repair so they are not polluting extra NO2. 2015 International Residential Code that requires that vent hoods are used for all stoves and set standards for residential buildings. This requires that all range hoods have a vent that discharges outside. You can also prevent NO2 exposure by avoiding cigarette smoking and not idling your car whenever possible.[37]

Environmental limits

The U.S. EPA has set safety levels for environmental exposure to NO
2
at 100 ppb, averaged over one hour, and 53 ppb, averaged annually.[8] As of February 2016, no area of the US was out of compliance with these limits and concentrations ranged between 10–20 ppb, and annual average ambient NO2 concentrations, as measured at area-wide monitors, have decreased by more than 40% since 1980.[33]

However, NO
2
concentrations in vehicles and near roadways are appreciably higher than those measured at monitors in the current network. In fact, in-vehicle concentrations can be 2–3 times higher than measured at nearby area-wide monitors. Near-roadway (within about 50 metres (160 ft)) concentrations of NO2 have been measured to be approximately 30 to 100% higher than concentrations away from roadways. Individuals who spend time on or near major roadways can experience short-term NO2 exposures considerably higher than measured by the current network. Approximately 16% of U.S. housing units are located within 300 feet (91 m) of a major highway, railroad, or airport (approximately 48 million people). Studies show a connection between breathing elevated short-term NO2 concentrations, and increased visits to emergency departments and hospital admissions for respiratory issues, especially asthma. NO2 exposure concentrations near roadways are of particular concern for susceptible individuals, including asthmatics, children, and the elderly.[29]

For limits in other countries see the table in the Ambient air quality criteria article.

See also

References

  1. ^ "nitrogen dioxide (CHEBI:33101)". Chemical Entities of Biological Interest (ChEBI). UK: European Bioinformatics Institute. 13 January 2008. Main. Retrieved 4 October 2011.
  2. ^ a b c d e Haynes, 4.79
  3. ^ Mendiara, S. N.; Sagedahl, A.; Perissinotti, L. J. (2001). "An electron paramagnetic resonance study of nitrogen dioxide dissolved in water, carbon tetrachloride and some organic compounds". Applied Magnetic Resonance. 20: 275–287. doi:10.1007/BF03162326.
  4. ^ Haynes, 4.134
  5. ^ Haynes, 5.16
  6. ^ a b c d e NIOSH Pocket Guide to Chemical Hazards. "#0454". National Institute for Occupational Safety and Health (NIOSH).
  7. ^ a b "Nitrogen dioxide". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  8. ^ a b c  This article incorporates public domain material from the United States Environmental Protection Agency document: "Nitrogen dioxide". United States Environmental Protection Agency. Feb 23, 2016.
  9. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 455. ISBN 978-0-08-037941-8.
  10. ^ a b c Holleman, A. F.; Wiberg, E. (2001) Inorganic Chemistry. Academic Press: San Diego. ISBN 0-12-352651-5.
  11. ^ Thiemann, Michael; Scheibler, Erich and Wiegand, Karl Wilhelm (2005). "Nitric Acid, Nitrous Acid, and Nitrogen Oxides". Ullmann's Encyclopedia of Industrial Chemistry. Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a17_293. ISBN 978-3527306732.CS1 maint: Uses authors parameter (link)
  12. ^ Finlayson-Pitts, B. J.; Wingen, L. M.; Sumner, A. L.; Syomin, D.; Ramazan, K. A. (2002-12-16). "The heterogeneous hydrolysis of NO2 in laboratory systems and in outdoor and indoor atmospheres: An integrated mechanism". Physical Chemistry Chemical Physics. 5 (2): 223–242. doi:10.1039/B208564J.
  13. ^ WHO Air Quality Guidelines – Second Edition. Chapter 7.1 Nitrogen Dioxide
  14. ^ Subcommittee on Emergency and Continuous Exposure Guidance Levels for Selected Submarine Contaminants; Committee on Toxicology; Board on Environmental Studies and Toxicology; Division on Earth and Life Studies; National Research Council. Chapter 12: Nitrogen Dioxide in Emergency and Continuous Exposure Guidance Levels for Selected Submarine Contaminants. National Academies Press, 2007. ISBN 978-0-309-09225-8
  15. ^ "Mechanism Overview, June 2012" (PDF). noxilizer.com. Noxilizer, Inc. Retrieved 2 July 2013.
  16. ^ a b Cotton, Simon (21 March 2013) Nitrogen dioxide. RSC Chemistry World
  17. ^ "WHO | Air quality guidelines – global update 2005". www.who.int. Retrieved 2016-10-19.
  18. ^ US Dept. of Health and Human Services, Public Health Service, Agency for Toxic Substances and Disease Registry, Division of Toxicology. April 2002 ATSDR Nitrous Oxides
  19. ^ "The Impact of Unvented Gas Heating Appliances on Indoor Nitrogen Dioxide Levels in 'TIGHT' Homes" (PDF). ahrinet.org. 2013-03-21.
  20. ^ Chan-Yeung, M.; Ashley, M. J.; Grzybowski, S. (1978). "Grain dust and the lungs". Canadian Medical Association Journal. 118 (10): 1271–4. PMC 1818652. PMID 348288.
  21. ^ Gurney, J. W.; Unger, J. M.; Dorby, C. A.; Mitby, J. K.; von Essen, S. G. (1991). "Agricultural disorders of the lung". Radiographics. 11 (4): 625–34. doi:10.1148/radiographics.11.4.1887117. PMID 1887117.
  22. ^ Effects of Nuclear Explosions. Nuclearweaponarchive.org. Retrieved on 2010-02-08.
  23. ^ a b U.S. EPA. Integrated Science Assessment for Oxides of Nitrogen – Health Criteria (2016 Final Report). U.S. Environmental Protection Agency, Washington, DC, EPA/600/R-15/068, 2016. Federal Register Notice Jan 28, 2016 Free download available at Report page at EPA website.
  24. ^ Toxnet Nitrogen dioxide: Human Health Effects Page accessed March 28, 2016
  25. ^ CDC NIOSH International Chemical Safety Cards (ICSC): Nitrogen Dioxide Page last reviewed: July 22, 2015; Page last updated: July 1, 2014
  26. ^ Agency for Toxic Substances and Disease Registry via the CDC Medical Management Guidelines for Nitrogen Oxides Page last reviewed: October 21, 2014; Page last updated: October 21, 2014
  27. ^ University of Kansas Hospital, Poison Control Center Poison Facts: Medium Chemicals: Nitrogen Dioxide page accessed March 28, 2016
  28. ^ "40 C.F.R.: Appendix A to Part 355—The List of Extremely Hazardous Substances and Their Threshold Planning Quantities" (PDF) (July 1, 2008 ed.). Government Printing Office. Retrieved October 29, 2011.
  29. ^ a b  This article incorporates public domain material from the United States Environmental Protection Agency document "Nitrogen Dioxide: Health". Retrieved on 2/23/2016.
  30. ^ a b c d Heinrich, Joachim (2011-01-01). "Influence of indoor factors in dwellings on the development of childhood asthma". International Journal of Hygiene and Environmental Health. 214 (1): 1–25. doi:10.1016/j.ijheh.2010.08.009. PMID 20851050.
  31. ^ a b Garrett, Maria H.; Hooper, Martin A.; Hooper, Beverley M.; Abramson, Michael J. (1998-09-01). "Respiratory Symptoms in Children and Indoor Exposure to Nitrogen Dioxide and Gas Stoves". American Journal of Respiratory and Critical Care Medicine. 158 (3): 891–895. doi:10.1164/ajrccm.158.3.9701084. PMID 9731022.
  32. ^ "Historical Census of Housing Tables -House Heating Fuel". www.census.gov. Retrieved 2016-10-19.
  33. ^ a b  This article incorporates public domain material from the United States Environmental Protection Agency document "Nitrogen Dioxide Basic Information". Retrieved on 2/23/2016.
  34. ^ Chapman, Robert S.; Hadden, Wilbur C.; Perlin, Susan A. (2003-07-15). "Influences of asthma and household environment on lung function in children and adolescents: the third national health and nutrition examination survey". American Journal of Epidemiology. 158 (2): 175–189. PMID 12851231.
  35. ^ Hansel, Nadia N.; Breysse, Patrick N.; McCormack, Meredith C.; Matsui, Elizabeth C.; Curtin-Brosnan, Jean; Williams, D’Ann L.; Moore, Jennifer L.; Cuhran, Jennifer L.; Diette, Gregory B. (2016-10-19). "A Longitudinal Study of Indoor Nitrogen Dioxide Levels and Respiratory Symptoms in Inner-City Children with Asthma". Environmental Health Perspectives. 116 (10): 1428–1432. doi:10.1289/ehp.11349. PMC 2569107. PMID 18941590.
  36. ^ Kile, Molly L; Coker, Eric S; Smit, Ellen; Sudakin, Daniel; Molitor, John; Harding, Anna K (2014-09-02). "A cross-sectional study of the association between ventilation of gas stoves and chronic respiratory illness in U.S. children enrolled in NHANESIII". Environmental Health. 13: 71. doi:10.1186/1476-069X-13-71. PMC 4175218. PMID 25182545.
  37. ^ "Healthy Child Healthy World". Healthy Child Healthy World. Archived from the original on 2016-10-11. Retrieved 2016-10-19.

Cited sources

External links

Bent molecular geometry

In chemistry, the term "bent" can be applied to certain molecules to describe their molecular geometry. Certain atoms, such as oxygen, will almost always set their two (or more) covalent bonds in non-collinear directions due to their electron configuration. Water (H2O) is an example of a bent molecule, as well as its analogues. The bond angle between the two hydrogen atoms is approximately 104.45°. Nonlinear geometry is commonly observed for other triatomic molecules and ions containing only main group elements, prominent examples being nitrogen dioxide (NO2), sulfur dichloride (SCl2), and methylene (CH2).

This geometry is almost always consistent with VSEPR theory, which usually explains non-collinearity of atoms with a presence of lone pairs. There are several variants of bending, where the most common is AX2E2 where two covalent bonds and two lone pairs of the central atom (A) form a complete 8-electron shell. They have central angles from 104° to 109.5°, where the latter is consistent with a simplistic theory which predicts the tetrahedral symmetry of four sp3 hybridised orbitals. The most common actual angles are 105°, 107°, and 109°: they vary because of the different properties of the peripheral atoms (X).

Other cases also experience orbital hybridisation, but in different degrees. AX2E1 molecules, such as SnCl2, have only one lone pair and the central angle about 120° (the centre and two vertices of an equilateral triangle). They have three sp2 orbitals. There exist also sd-hybridised AX2 compounds of transition metals without lone pairs: they have the central angle about 90° and are also classified as bent.

Dinitrogen tetroxide

Dinitrogen tetroxide, commonly referred to as nitrogen tetroxide, is the chemical compound N2O4. It is a useful reagent in chemical synthesis. It forms an equilibrium mixture with nitrogen dioxide.

Dinitrogen tetroxide is a powerful oxidizer that is hypergolic (spontaneously reacts) upon contact with various forms of hydrazine, which has made the pair a common bipropellant for rockets.

Dinitrogen trioxide

Dinitrogen trioxide is the chemical compound with the formula N2O3. This deep blue solid is one of the simple nitrogen oxides. It forms upon mixing equal parts of nitric oxide and nitrogen dioxide and cooling the mixture below −21 °C (−6 °F):

NO + NO2 ⇌ N2O3Dinitrogen trioxide is only isolable at low temperatures, i.e. in the liquid and solid phases. At higher temperatures the equilibrium favors the constituent gases, with Kdiss = 193 kPa (25 °C).

Lead chamber process

The lead chamber process was an industrial method used to produce sulfuric acid in large quantities. It has been largely supplanted by the contact process.

In 1746 in Birmingham, England, John Roebuck began producing sulfuric acid in lead-lined chambers, which were stronger and less expensive, and could be made much larger, than the glass containers which had been used previously. This allowed the effective industrialization of sulfuric acid production and, with several refinements, this process remained the standard method of production for almost two centuries. So robust was the process that as late as 1946, the chamber process still accounted for 25% of sulfuric acid manufactured.

Leighton relationship

In atmospheric chemistry, the Leighton relationship is an equation that determines the concentration of tropospheric ozone in areas polluted by the presence of nitrogen oxides. Ozone in the troposphere is primarily produced through the photolysis of nitrogen dioxide by photons with wavelengths (λ) less than 420 nanometers, which are able to reach the lowest levels of the atmosphere, through the following mechanism:

NO2 + hν (λ < 420 nm) → NO + O (3P)

 

 

 

 

(J1)

O (3P) + O2 + M → O3 + M

 

 

 

 

(k2)

NO + O3 → NO2 + O2

 

 

 

 

(k3)

This series of reactions creates a null cycle, in which there is no net production or loss of any species involved. Since O (3P) is very reactive and O2 is abundant, O (3P) can be assumed to be in steady state, and thus an equation linking the concentrations of the species involved can be derived:

The Leighton relationship above shows how production of ozone is directly related to the solar intensity, and hence to the zenith angle, due to the reliance on photolysis of NO2. The yield of ozone will therefore be greatest during the day, especially at noon and during the summer season. This relationship also demonstrates how high concentrations of both ozone and nitric oxide are unfeasible. However, NO can react with peroxyl radicals to produce NO2 without loss of ozone:

RO2 + NO → NO2 + RO

thus providing another pathway to allow for the buildup of ozone by breaking the above null cycle.

This relationship is named after Philip Leighton, author of the groundbreaking 1961 book Photochemistry of Air Pollution, as recognition of his contributions in the understanding of tropospheric chemistry. Computer models of atmospheric chemistry utilize the Leighton relationship to minimize complexity by deducing the concentration of one of ozone, nitrogen dioxide, and nitric oxide when the concentrations of the other two are known.

Mixed oxides of nitrogen

Mixed oxides of nitrogen (MON) are solutions of nitric oxide (NO) in dinitrogen tetroxide/nitrogen dioxide (N2O4 and NO2). It may be used as an oxidizing agent in rocket propulsion systems. A broad range of compositions is available, and can be denoted as MONi, where i represents the percentage of nitric oxide in the mixture (e.g. MON3 contains 3% nitric oxide, MON25 25% nitric oxide). An upper limit is MON40 (40% by weight). In Europe MON 1.3 is mostly used for rocket propulsion systems, while the NASA seems to prefer MON 3. A higher percentage of NO decreases the corrosiveness of the liquid, while the costs increase and the oxidation potential is decreased.

The addition of nitric oxide also reduces the freezing point to a more desirable temperature. The freezing point of pure nitrogen tetroxide is −9 °C (16 °F), while MON3 is −15 °C (5 °F) and MON25 is −55 °C (−67 °F).

NOx Law (Japan)

Amendment Act on Reduction of Total Amount of Nitrogen Dioxide and Particulate Matters Originating from Automobiles in Designated Areas (introduced on 3 June 1992, amended on 19 June 2001) is environmental legislation in Japan.

Nitric acid

Nitric acid (HNO3), also known as aqua fortis (Latin for "strong water") and spirit of niter, is a highly corrosive mineral acid.

The pure compound is colorless, but older samples tend to acquire a yellow cast due to decomposition into oxides of nitrogen and water. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86% HNO3, it is referred to as fuming nitric acid. Depending on the amount of nitrogen dioxide present, fuming nitric acid is further characterized as white fuming nitric acid at concentrations above 95%, or red fuming nitric acid at concentrations above 86%.

Nitric acid is the primary reagent used for nitration – the addition of a nitro group, typically to an organic molecule. While some resulting nitro compounds are shock- and thermally-sensitive explosives, a few are stable enough to be used in munitions and demolition, while others are still more stable and used as pigments in inks and dyes. Nitric acid is also commonly used as a strong oxidizing agent.

Nitric oxide

Nitric oxide (nitrogen oxide 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.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. It was proclaimed the "Molecule of the Year" in 1992. 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.

Nitrogen dioxide poisoning

Nitrogen dioxide poisoning is the illness resulting from the toxic effect of nitrogen dioxide (NO2). It usually occurs after the inhalation of the gas beyond the threshold limit value.

Nitrogen dioxide is reddish-brown with very harsh smell at high concentrations. It is colorless and odorless at lower concentration but still harmful. Nitrogen dioxide poisoning depends on the duration, frequency, and intensity of exposure.

Nitrogen dioxide is an irritant of the mucous membrane linked with another air pollutant that causes pulmonary diseases such as OLD, asthma, chronic obstructive pulmonary disease and sometimes acute exacerbation of COPD and in fatal cases, deaths.

Its poor solubility in water enhances its passage and its ability to pass through the moist oral mucosa of the respiratory tract.

Like most toxic gases, the dose inhaled determines the toxicity on the respiratory tract. Occupational exposures constitute the highest risk of toxicity and domestic exposure is uncommon. Prolonged exposure to low concentration of the gas may have lethal effects, as can short-term exposure to high concentrations like chlorine gas poisoning. It is one of the major air pollutant capable of causing severe heath hazards such as coronary artery disease as well as stroke.

Nitrogen dioxide is often released into the environment as a byproduct of fuel combustion but rarely released by spontaneous combustion. Known sources of nitrogen dioxide gas poisoning include automobile exhaust and power stations.

The toxicity may also result from non-combustible sources such as the one released from anaerobic fermentation of food grains and anaerobic digestion of biodegradable waste.

The World Health Organization (WHO) developed a global recommendation limiting exposures to less than 20 parts per billion for chronic exposure and value less 100 ppb for one hour for acute exposure, using nitrogen dioxide as a marker for other pollutants from fuel combustion.

The standards also based on the concentration of nitrogen dioxide that show a significant and profound effects on the function of the pulmonary of asthmatic patients.

Historically, some cities in the United States including Chicago and Los Angeles have high levels of nitrogen dioxide but the EPA set a standard values less than 100 ppb for a one-hour exposure and less than 53 ppb for chronic exposure.

Nitronium ion

The nitronium ion, NO+2, is a cation. It is an onium ion because of its tetravalent nitrogen atom and +1 charge, similar in that regard to ammonium. It is created by the removal of an electron from the paramagnetic nitrogen dioxide molecule, or the protonation of nitric acid (with removal of H2O).

It is stable enough to exist in normal conditions, but it is generally reactive and used extensively as an electrophile in the nitration of other substances. The ion is generated in situ for this purpose by mixing concentrated sulfuric acid and concentrated nitric acid according to the equilibrium:

H2SO4 + HNO3 → HSO−4 + NO+2 + H2O

Nitryl

Nitryl is the nitrogen dioxide (NO2) moiety when it occurs in a larger compound as a univalent fragment. Examples include nitryl fluoride (NO2F) and nitryl chloride (NO2Cl).Like nitrogen dioxide, the nitryl moiety contains a nitrogen atom with two bonds to the two oxygen atoms, and a third bond shared equally between the nitrogen and the two oxygen atoms. The nitrogen-centred radical is then free to form a bond with another univalent fragment (X) to produce an N-X bond, where X can be F, Cl, OH, etc.

In organic nomenclature, the nitryl moiety is known as the nitro group. For instance, nitryl benzene is normally called nitrobenzene (PhNO2).

Oxidized cellulose

Oxidized cellulose is a water-insoluble derivative of cellulose. It can be produced from cellulose by the action of an oxidizing agent, such as chlorine, hydrogen peroxide, peracetic acid, chlorine dioxide, nitrogen dioxide, persulfates, permanganate, dichromate-sulfuric acid, hypochlorous acid, hypohalites or periodates and a variety of metal catalysts. Oxidized cellulose may contain carboxylic acid, aldehyde, and/or ketone groups, in addition to the original hydroxyl groups of the starting material, cellulose, depending on the nature of the oxidant and reaction conditions. It is an antihemorrhagic.

Photoinitiator

A photoinitiator is a molecule that creates reactive species (free radicals, cations or anions) when exposed to radiation (UV or visible). Synthetic photoinitiators are key components in photopolymers (i.e., photo-curable coatings, adhesives and dental restoratives).

Some small molecules in the atmosphere can also act as photoinitiators by decomposing to give free radicals (in photochemical smog). For instance, nitrogen dioxide is produced in large quantities by gasoline-burning internal combustion engines. NO2 in the troposphere gives smog its brown coloration and catalyzes production of toxic ground-level ozone. Molecular oxygen (O2) also serves as a photoinitiator in the stratosphere, breaking down into atomic oxygen and combining with O2 in order to form the ozone in the ozone layer.

Reactive nitrogen species

Reactive nitrogen species (RNS) are a family of antimicrobial molecules derived from nitric oxide (•NO) and superoxide (O2•−) produced via the enzymatic activity of inducible nitric oxide synthase 2 (NOS2) and NADPH oxidase respectively. NOS2 is expressed primarily in macrophages after induction by cytokines and microbial products, notably interferon-gamma (IFN-γ) and lipopolysaccharide (LPS).Reactive nitrogen species act together with reactive oxygen species (ROS) to damage cells, causing nitrosative stress. Therefore, these two species are often collectively referred to as ROS/RNS.

Reactive nitrogen species are also continuously produced in plants as by-products of aerobic metabolism or in response to stress.

Red fuming nitric acid

Red fuming nitric acid (RFNA) is a storable oxidizer used as a rocket propellant. It consists of 84% nitric acid (HNO3), 13% dinitrogen tetroxide and 1–2% water. The color of red fuming nitric acid is due to the dinitrogen tetroxide, which breaks down partially to form nitrogen dioxide. The nitrogen dioxide dissolves until the liquid is saturated, and evaporates off into fumes with a suffocating odor. RFNA increases the flammability of combustible materials and is highly exothermic when reacting with water.

It is usually used with an inhibitor (with various, sometimes secret, substances, including hydrogen fluoride; any such combination is called inhibited RFNA, IRFNA) because nitric acid attacks most container materials.

It can also be a component of a monopropellant; with substances like amine nitrates dissolved in it, it can be used as the sole fuel in a rocket. It is not normally used this way however.

During World War II, the German military used RFNA in some rockets. The mixtures used were called S-Stoff (96% nitric acid with 4% ferric chloride as an ignition catalyst) and SV-Stoff (94% nitric acid with 6% dinitrogen tetroxide) and nicknamed Salbei (sage).

Inhibited RFNA was the oxidizer of the world's most-launched light orbital rocket, the Kosmos-3M.

Other uses for RFNA include fertilizers, dye intermediates, explosives, and pharmaceutic aid as acidifier. It can also be used as a laboratory reagent in photoengraving and metal etching.

SAGE III on ISS

SAGE III on ISS is the fourth generation of a series of NASA Earth-observing instruments, known as the Stratospheric Aerosol and Gas Experiment. The first SAGE III instrument was launched on the Russian Meteor (satellite) spacecraft. The recently revised SAGE III will be mounted to the International Space Station where it will use the unique vantage point of ISS to make long-term measurements of ozone, aerosols, water vapor, and other gases in Earth's atmosphere.

Solar Mesosphere Explorer

The Solar Mesosphere Explorer (also known as Explorer 64) was a United States unmanned spacecraft to investigate the processes that create and destroy ozone in Earth's upper atmosphere. The mesosphere is a layer of the atmosphere extending from the top of the stratosphere to an altitude of about 80 kilometers (50 mi). The spacecraft carried five instruments to measure ozone, water vapor and incoming solar radiation.

Launched on October 6, 1981, on a Delta rocket from Vandenberg Air Force Base, in California, the satellite returned data until April 4, 1989. The spacecraft reentered Earth's atmosphere on March 5, 1991.

Managed for NASA by the Jet Propulsion Laboratory, the Solar Mesosphere Explorer was built by Ball Space Systems and operated by the Laboratory for Atmospheric and Space Physics of the University of Colorado where one hundred undergraduate and graduate students were involved.

Mass: 437 kilograms (963 pounds)

Power: Solar panels which charged NiCad batteries

Configuration: Cylinder 1.25 meter (4.1 ft) diameter by 1.7 meter (5.6 ft) high

Science instruments: Ultraviolet ozone spectrometer, 1.27 micrometre spectrometer, nitrogen dioxide spectrometer, four-channel infrared radiometer, solar ultraviolet monitor, solar proton alarm detector

Nitrogen species
Mixed oxidation states
+1 oxidation state
+2 oxidation state
+3 oxidation state
+4 oxidation state
+5 oxidation state
+6 oxidation state
+7 oxidation state
+8 oxidation state
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