|Jmol-3D images||Image 1|
|Molar mass||92.011 g/mol|
|Appearance||colourless liquid / orange gas|
|Density||1.44246 g/cm3 (liquid, 21 °C)|
−11.2 °C, 262 K, 12 °F
21.69 °C, 295 K, 71 °F
|Solubility in water||reacts|
|Vapor pressure||96 kPa (20 °C)|
|Refractive index (nD)||1.00112|
|Molecular shape||planar, D2h|
|Std enthalpy of
|304.29 J K−1 mol−1|
|EU classification||Very toxic (T+)
|S-phrases||(S1/2), S9, S26, S28, S36/37/39, S45|
|Related nitrogen oxides||Nitrous oxide
| (what is: / ?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Dinitrogen tetroxide (nitrogen tetroxide) is the chemical compound N2O4. It is a useful reagent in chemical synthesis. It forms an equilibrium mixture with nitrogen dioxide; some call this mixture dinitrogen tetroxide, while some call it nitrogen dioxide.
Dinitrogen tetroxide is a powerful oxidizer. N2O4 is hypergolic with various forms of hydrazine, i.e., they burn on contact without a separate ignition source, making them popular bipropellant rocket fuels.
Structure and properties
Dinitrogen tetroxide forms an equilibrium mixture with nitrogen dioxide. 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.
Unlike NO2, N2O4 is diamagnetic since it has no unpaired electrons. 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.
Use as a rocket propellant
Dinitrogen 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, Apollo spacecraft and 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. It is also the primary oxidizer for Russia's Proton rocket and China's Long March rockets.
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). 
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 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 crewmember lost consciousness during descent. Upon landing, the crew was hospitalized 14 days for chemical-induced pneumonia and edema.
Power generation using N2O4
|This section needs additional citations for verification. (May 2013)|
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. 
Intermediate in the manufacture of nitric acid
- N2O4 + H2O → HNO2 + HNO3
- 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
N2O4 behaves as the salt [NO+] [NO3−], the former being a strong oxidant:
- 2 N2O4 + M → 2 NO + M(NO3)2
If metal nitrates are prepared from NTO 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 degrees 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.
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