Ethyl nitrate

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Ethyl nitrate
Skeletal formula of ethyl nitrate
Ball-and-stick model of the ethyl nitrate molecule
IUPAC name
Ethyl nitrate
Other names
Nitric acid ethyl ester
3D model (JSmol)
ECHA InfoCard 100.009.913
Molar mass 91.07 g/mol
Appearance colorless liquid
Density 1.10g/cm3
Melting point −102 °C (−152 °F; 171 K)
Boiling point 87.5 °C (189.5 °F; 360.6 K)
NFPA 704
Flammability code 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g., gasolineHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroformReactivity code 4: Readily capable of detonation or explosive decomposition at normal temperatures and pressures. E.g., nitroglycerinSpecial hazards (white): no codeNFPA 704 four-colored diamond
Flash point −37 °C; −34 °F; 236 K
Explosive limits 4.1%-50%
Related compounds
Methyl nitrate
Ethylene glycol dinitrate
Isopropyl nitrate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Ethyl nitrate is the ethyl ester of nitric acid and has the chemical formula C2H5NO3. It is a colourless, volatile, highly flammable liquid. It is used in organic synthesis and as an intermediate in the preparation of some drugs, dyes, and perfumes.[1]

Ethyl nitrate is found in the atmosphere, where it can react with other gases to form smog. Originally thought to be a pollutant, formed mainly by the combustion of fossil fuels, recent analysis of ocean water samples reveal that in places where cool water rises from the deep, the water is saturated with alkyl nitrates, likely formed by natural processes.[2]


Ethyl nitrate has been prepared by bubbling gaseous nitryl fluoride through ethanol at −10 °C.[3] The reaction was subsequently studied in detail.[4][5]

Ethyl nitrate can be prepared by nitrating ethanol with fuming nitric acid or a mixture of concentrated sulfuric and nitric acids. Further purifying by distillation carries a risk of explosion.[6]


  1. ^ 1921-, Schofield, Kenneth, (1980). Aromatic nitration. Cambridge: Cambridge University Press. p. 94. ISBN 9780521233620. OCLC 6357479.
  2. ^ S. Perkins (August 12, 2002). "Ocean yields gases that had seemed humanmade". Science News.
  3. ^ G. Hetherington and R. L. Robinson (1954). "Nitryl fluoride as a nitrating agent". J. Chem. Soc.: 3512. doi:10.1039/JR9540003512.
  4. ^ B. S. Fedorov and L. T. Eremenko (1997). "Nitration of alcohols by nitryl fluoride". Russian Chemical Bulletin. 46 (5): 1022–1023. doi:10.1007/BF02496138.
  5. ^ Explosives, 6th Edition, R. Meyer, J. Kohler, A. Homburg; page 125
  6. ^ Cohen, Julius B. (Julius Berend) (1920). Theoretical organic chemistry. University of California Libraries. London, Macmillan. p. 189.