Ammonium carbamate

Ammonium carbamate
Names
Other names
ammonium aminoformate, carbamic acid ammoniate, carbamic acid ammonium salt, carbamic acid monoammonium salt[1]
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.012.896
EC Number 214-185-2
14637 (G)
RTECS number EY8575000
UN number 9083
Properties
CH6N2O2
Molar mass 78.07 g·mol−1
Appearance Colorless, rhombic crystals
Density 1.38 g/cm3 (20 °C)
Melting point 152 °C (306 °F; 425 K)
Boiling point 251 °C (484 °F; 524 K) at 760 mmHg
Freely soluble in water
Solubility Soluble in alcohol
log P −0.47 in octanol/water
Vapor pressure 492 mmHg(51 °C)
Thermochemistry
151.8 kj/mol
Hazards
Main hazards Harmful if ingested, harmful to aquatic life, harmful if inhaled, respiatory tract irritation, skin irritation, eye irritation
Safety data sheet External MSDS
GHS pictograms
GHS signal word WARNING
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
0
1
0
Flash point 105.6 °C (222.1 °F; 378.8 K)
Lethal dose or concentration (LD, LC):
1,470 mg/kg in a rat
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Ammonium carbamate is the inorganic compound with the formula NH4[H2NCO2]. This salt formed by the reaction of ammonia with carbon dioxide. This compound is a white solid that is extremely soluble in water, less so in alcohol. It is unusual in degrading at room temperature.

Preparation and structure

It is prepared by the direct reaction between liquid ammonia and dry ice (solid carbon dioxide):[2]

2 NH3 + CO2 → H2NCOONH4

The structure of solid ammonium carbamate has been confirmed by X-ray crystallography. The oxygen centers form hydrogen bonds to the ammonium cation.[3]

Reactions

Ammonium carbamate reverts to carbon dioxide and ammonia even as a solid:[4]

NH2CO2NH4 → 2NH3 + CO2

It hydrates reversibly:[4]

NH2CO2NH4 + H2O → (NH4)2CO3

Ammonium carbamate serves a key role in the formation of carbamoyl phosphate, which is necessary for both the urea cycle and the production of pyrimidines. In this enzyme-catalyzed reaction, ATP and ammonium carbamate are converted to ADP and carbamoyl phosphate:[5][6]

ATP + NH2CO2NH4 → ADP + H2NC(O)OPO32−

Uses

The ability of ammonium carbamate to make urea was first discovered in 1870 when Bassarov heated ammonium carbamate in sealed glass tubes at temperatures ranging from 130 to 140 °C. This heating yields urea and water in an equimolar ratio.[4] A typical industrial plant that makes urea can produce up to 1500 tons a day. Ammonia and carbon dioxide is excessively fed to a synthesis reactor in this process. Ammonium carbamate is produced as an intermediate in this reactor and can then be dehydrated to urea according to the following equation:[7]

NH2CO2NH4 → NH2CONH2 + H2O

Ammonium carbamate has also been approved by the Environmental Protection Agency as an inert ingredient present in aluminum phosphide pesticide formulations. This pesticide is commonly used for insect and rodent control in areas where agricultural products are stored. The reason for ammonium carbamate as an ingredient is to make the phosphine less flammable by freeing ammonia and carbon dioxide to dilute hydrogen phosphide formed by a hydrolysis reaction.[8]

Ammonium carbamate can also be used as a good ammoniating agent. Compared to ammonia however, it is not nearly as strong an ammoniating agent.[9]

References

  1. ↑ "Ammonium Carbamate" Retrieved October 12, 2012.
  2. ↑ Brooks, L. A.; Audrieta, L. F.; Bluestone, H.; Jofinsox, W. C. (2007). "Ammonium Carbamate". Inorg. Synth. 2: 85. doi:10.1002/9780470132333.ch23.
  3. ↑ J. M. Adams; R. W. H. Small (1973). "The crystal structure of ammonium carbamate". Acta Cryst. B29: 2317–2319. doi:10.1107/S056774087300662X.
  4. 1 2 3 Clark, K. G.; Gaddy, V. L.; Rist, C. E. (1933). "Equilibria in the Ammonium Carbamate-Urea-Water System". Ind. Eng. Chem. 25 (10): 1092–1096. doi:10.1021/ie50286a008.
  5. ↑ Goldberg, R. N. Apparent Equilibrium Constants for Enzyme-catalyzed reactions (2009). CRC Handbook of Chemistry and Physics, 7–19. Retrieved from https://www.nist.gov/manuscript-publication-search.cfm?pub_id=900943
  6. ↑ Phosphorus Compounds: Advances in Research and Application: 2011 Edition
  7. ↑ Barzagli, F., Mani, F., & Peruzzini, M. (2011). From greenhouse gas to feedstock: formation of ammonium carbamate from CO2 and NH3 in organic solvents and its catalytic conversion into urea under mild conditions. Green Chemistry, 13(5), 1267–1274
  8. ↑ United States Environmental Protection Agency. (2006). Inert Reassessment-Ammonium Carbamate [Data File]. Retrieved from http://www.epa.gov/opprd001/inerts/carbamate.pdf
  9. ↑ "Merck Index Fourteenth Addition" Retrieved October 19, 2012.
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