Dinitrogen tetroxide

From Wikipedia, the free encyclopedia
  (Redirected from Nitrogen tetroxide)
Jump to: navigation, search
Dinitrogen tetroxide
Full structural formula
Space-filling model
Nitrogen dioxide at different temperatures
IUPAC name
Dinitrogen tetroxide
Other names
Dinitrogen(II) oxide(-I)
10544-72-6 YesY
ChEBI CHEBI:29803 YesY
ChemSpider 23681 YesY
EC number 234-126-4
Jmol-3D images Image
PubChem 25352
RTECS number QW9800000
UN number 1067
Molar mass 92.011 g/mol
Appearance colourless liquid / orange gas
Density 1.44246 g/cm3 (liquid, 21 °C)
Melting point −11.2 °C (11.8 °F; 261.9 K)
Boiling point 21.69 °C (71.04 °F; 294.84 K)
Vapor pressure 96 kPa (20 °C)[1]
Molecular shape planar, D2h
Dipole moment zero
304.29 J K−1 mol−1[2]
+9.16 kJ/mol[2]
MSDS External MSDS
EU Index 007-002-00-0
EU classification Very toxic (T+)
Corrosive (C)
R-phrases R26, R34
S-phrases (S1/2), S9, S26, S28, S36/37/39, S45
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Flash point Non-flammable
Related compounds
Nitrous oxide
Nitric oxide
Dinitrogen trioxide
Nitrogen dioxide
Dinitrogen pentoxide
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 YesY verify (what isYesY/N?)
Infobox references

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 makes the pair a popular bipropellant for rockets.

Structure and properties[edit]

Dinitrogen tetroxide forms an equilibrium mixture with nitrogen dioxide.[3] The molecule is planar with an N-N bond distance of 1.78 Å and N-O distances of 1.19 Å. The N-N distance corresponds to a weak bond, since it is significantly longer than the average N-N single bond length of 1.45 Å.[4]

Unlike NO2, N2O4 is diamagnetic since it has no unpaired electrons.[5] The liquid is also colorless but can appear as a brownish yellow liquid due to the presence of NO2 according to the following equilibrium:

N2O4 2 NO2

Higher temperatures push the equilibrium towards nitrogen dioxide. Inevitably, some dinitrogen tetroxide is a component of smog containing nitrogen dioxide.


Nitrogen dioxide is made by the catalytic oxidation of ammonia: steam is used as a diluent to reduce the combustion temperature. Most of the water is condensed out, and the gases are further cooled; the nitric oxide that was produced is oxidized to nitrogen dioxide, and the remainder of the water is removed as nitric acid. The gas is essentially pure nitrogen tetroxide, which is condensed in a brine-cooled liquefier.[citation needed]

Use as a rocket propellant[edit]

Nitrogen tetroxide is one of the most important rocket propellants ever developed, much like the German-developed hydrogen peroxide–based T-Stoff oxidizer used in their World War II rocket-propelled combat aircraft designs such as the Messerschmitt Me 163 Komet, and by the late 1950s it became the storable oxidizer of choice for rockets in both the USA and USSR. It is a hypergolic propellant often used in combination with a hydrazine-based rocket fuel. One of the earliest uses of this combination was on the Titan rockets used originally as ICBMs and then as launch vehicles for many spacecraft. Used on the U.S. Gemini and Apollo spacecraft and also on the Space Shuttle, it continues to be used on most geo-stationary satellites, and many deep-space probes. It now seems likely that NASA will continue to use this oxidizer in the next-generation 'crew-vehicles' which will replace the shuttle.[citation needed] It is also the primary oxidizer for Russia's Proton rocket.

When used as a propellant, dinitrogen tetroxide is usually referred to simply as 'Nitrogen Tetroxide' and the abbreviation 'NTO' is extensively used. Additionally, NTO is often used with the addition of a small percentage of nitric oxide, which inhibits stress-corrosion cracking of titanium alloys, and in this form, propellant-grade NTO is referred to as "Mixed Oxides of Nitrogen" or "MON". Most spacecraft now use MON instead of NTO; for example, the Space Shuttle reaction control system uses MON3 (NTO containing 3wt%NO).[6]

The Apollo-Soyuz mishap[edit]

On 24 July 1975, NTO poisoning affected the three U.S. astronauts on board the Apollo-Soyuz Test Project during its final descent. This was due to a switch negligently, or accidentally, left in the wrong position, which allowed NTO fumes to vent out of the Apollo spacecraft then back in through the cabin air intake from the outside air after the external vents were opened. One crew member lost consciousness during descent. Upon landing, the crew was hospitalized for 14 days for chemical-induced pneumonia and edema.[7]

Power generation using N2O4[edit]

The tendency of N2O4 to reversibly break into NO2 has led to research into its use in advanced power generation systems as a so-called dissociating gas. "Cool" nitrogen tetroxide is compressed and heated, causing it to dissociate into nitrogen dioxide at half the molecular weight. This hot nitrogen dioxide is expanded through a turbine, cooling it and lowering the pressure, and then cooled further in a heat sink, causing it to recombine into nitrogen tetroxide at the original molecular weight. It is then much easier to compress to start the entire cycle again. Such dissociative gas Brayton cycles have the potential to considerably increase efficiencies of power conversion equipment. [8]

Chemical reactions[edit]

Intermediate in the manufacture of nitric acid[edit]

Nitric acid is manufactured on a large scale via N2O4. This species reacts with water to give both nitrous acid and nitric acid:

N2O4 + H2O → HNO2 + HNO3

The coproduct HNO2 upon heating disproportionates to NO and more nitric acid. When exposed to oxygen, NO is converted back into nitrogen dioxide:

2 NO + O2 → 2 NO2

The resulting NO2 (and N2O4, obviously) can be returned to the cycle to give the mixture of nitrous and nitric acids again.

Synthesis of metal nitrates[edit]

N2O4 behaves as the salt [NO+] [NO3], the former being a strong oxidant:

2 N2O4 + M → 2 NO + M(NO3)2

where M = Cu, Zn, or Sn.

If metal nitrates are prepared from N2O4 in completely anhydrous conditions, a range of covalent metal nitrates can be formed with many transition metals. This is because there is a thermodynamic preference for the nitrate ion to bond covalently with such metals rather than form an ionic structure. Such compounds must be prepared in anhydrous conditions, since the nitrate ion is a much weaker ligand than water, and if water is present the simple hydrated nitrate will form. The anhydrous nitrates concerned are themselves covalent, and many, e.g. anhydrous copper nitrate, are volatile at room temperature. Anhydrous titanium nitrate sublimes in vacuum at only 40 °C. Many of the anhydrous transition metal nitrates have striking colours. This branch of chemistry was developed by Clifford Addisson and Noramn Logan at Nottingham University in the UK during the 1960s and 1970s when highly efficient desiccants and dry boxes started to become available.

See also[edit]


  1. ^ International Chemical Safety Card
  2. ^ a b P.W. Atkins and J. de Paula, Physical Chemistry (8th ed., W.H. Freeman, 2006) p.999
  3. ^ Henry A. Bent Dimers of Nitrogen Dioxide. II. Structure and Bonding Inorg. Chem., 1963, 2 (4), pp 747–752
  4. ^ R.H. Petrucci, W.S. Harwood and F.G. Herring General Chemistry (8th ed., Prentice-Hall 2002), p.420
  5. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  6. ^ [1]
  7. ^ Sotos, John G., MD. "Astronaut and Cosmonaut Medical Histories", May 12, 2008, accessed April 1, 2011.
  8. ^ Ragheb, R. "Nuclear Reactors Concepts and Thermodynamic Cycles". Retrieved 1 May 2013. 

External links[edit]