Ammonium carbamate

Ammonium carbamate is an inorganic compound with the formula NH4[H2NCO2] consisting of ammonium NH+
and carbamate H
. It is a white solid that is extremely soluble in water, less so in alcohol. Ammonium carbamate can be formed by the reaction of ammonia with carbon dioxide, and will slowly decompose to those gases at ordinary temperatures and pressures. It is an intermediary in the industrial synthesis of urea, an important fertilizer.

Ammonium carbamate
Ammonium carbamate 2-d.gif
Other names
ammonium amidocarbonate, azanium amidodioxidocarbonate, ammonium aminoformate, carbamic acid ammonium salt, carbamic acid monoammonium salt[1]
3D model (JSmol)
ECHA InfoCard 100.012.896 Edit this at Wikidata
EC Number
  • 214-185-2
14637 (G)
RTECS number
  • EY8575000
UN number 9083
Molar mass 78.071 g·mol−1
Appearance Colorless, rhombic crystals
Density 1.38 g/cm3 (20 °C)
Melting point 60 °C (140 °F; 333 K) decomposes
Freely soluble in water
Solubility Soluble in ethanol, methanol, liquid ammonia, formamide[2][3]
log P −0.47 in octanol/water
Vapor pressure 492 mmHg(51 °C)
-642.5 kJ/mol
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 GHS07: Harmful
GHS Signal word Warning
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
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


Solid-gas equilibriumEdit

In a closed container solid ammonium carbamate is in equilibrium with carbon dioxide and ammonia [4][5][6]

NH2CO2NH4 ⇌ 2NH3 + CO2

Lower temperatures shift the equilibrium towards the carbamate.

At higher temperatures ammonium carbamate condenses into urea:

+ H

This reaction was first discovered in 1870 by Bassarov, by heating ammonium carbamate in sealed glass tubes at temperatures ranging from 130 to 140 °C.[5]

Equilibrium in waterEdit

At ordinary temperatures and pressures, ammonium carbamate exists in aqueous solutions as an equilibrium with ammonia and carbon dioxide, and the anions bicarbonate HCO
and CO3,[7][5][8] Indeed, solutions of ammonium carbonate or bicarbonate will contain some carbamate anions too.

+ 2 H
+ HO
2 H
+ 2 H
⇌ 2 NH+
+ CO2−


The structure of solid ammonium carbamate has been confirmed by X-ray crystallography. The oxygen centers form hydrogen bonds to the ammonium cation.[9]

Natural occurrenceEdit

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:[10][11]



From liquid ammonia and dry iceEdit

Ammonium carbamate is prepared by the direct reaction between liquid ammonia and dry ice (solid carbon dioxide):[4]

2 NH3 + CO2 → H2NCOONH4

It is also prepared by reaction of the two gases at high temperature (450–500 K) and high pressure (150–250 bar).[12]

From gaseous ammonia and carbon dioxideEdit

Ammonium carbamate can also be obtained by bubbling gaseous CO
and NH
in anhydrous ethanol, 1-propanol, or DMF at ambient pressure and 0 °C. The carbamate precipitates and can be separated by simple filtration, and the liquid containing the unreacted ammonia can be returned to the reactor. The absence of water prevents the formation of bicarbonate and carbonate, and no ammonia is lost.[12]


Urea synthesisEdit

Ammonium carbamate is an intermediate in the industrial production of urea. A typical industrial plant that makes urea can produce up to 1500 tons a day. in this reactor and can then be dehydrated to urea according to the following equation:[12]


Pesticide formulationsEdit

Ammonium carbamate has also been approved by the Environmental Protection Agency as an inert ingredient present in aluminium 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 phosphine formed by a hydrolysis reaction.[13]


Ammonium carbamate can be used as a good ammoniating agent, though not nearly as strong as ammonia itself. For instance, it is an effective reagent for preparation of different substituted β-amino-α,β-unsaturated esters. The reaction can be carried out in methanol at room temperature and can be isolated in the absence of water, in high purity and yield.[14]

Preparation of metal carbamatesEdit

Ammonium carbamate can be a starting reagent for the production of salts of other cations. For instance, by reacting it with solid potassium chloride KCl in liquid ammonia one can obtain potassium carbamate K+
.[2] Carbamates of other metals, such as calcium, can be produced by reacting ammonium carbamate with a suitable salt of the desired cation, in an anhydrous solvent such as methanol, ethanol, or formamide, even at room temperature.[3]


  1. ^ "Ammonium Carbamate" Retrieved October 12, 2012.
  2. ^ a b Carl Theodor Thorssell and August Kristensson (1935): "Process for the production of potassium carbamate". US Patent 2002681, US31484228A
  3. ^ a b Erns Kuss and Emil Germann (1935): "Production of metal carbamates". US Patent US2023890A
  4. ^ a b Brooks, L. A.; Audrieta, L. F.; Bluestone, H.; Jofinsox, W. C. (1946). Ammonium Carbamate. Inorg. Synth. Inorganic Syntheses. 2. pp. 85–86. doi:10.1002/9780470132333.ch23. ISBN 9780470132333.
  5. ^ a b c 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.
  6. ^ R. N. Bennett, P. D. Ritchie, D. Roxburgh and J. Thomson (1953): "The system ammonia + carbon dioxide + ammonium carbamate. Part I. — The equilibrium of thermal dissociation of ammonium carbamate". Transactions of the Faraday Society, volume 49, pages 925-929. doi:10.1039/TF9534900925
  7. ^ George H. Burrows and Gilbert N. Lewis (1912): "The equilibrium between ammonium carbonate and ammonium carbamate in aqueous solution at 25°". Journal of the American Chemical Society, volume 34, issue 8, pages 993-995. doi:10.1021/ja02209a003
  8. ^ Fabrizio Mani, Maurizio Peruzzini, and Piero Stoppioni (2006): "CO2 absorption by aqueous NH
    solutions: speciation of ammonium carbamate, bicarbonate and carbonate by a 13C NMR study". Green Chemistry, volume 8, issue 11, pages 995-1000. doi:10.1039/B602051H
  9. ^ J. M. Adams; R. W. H. Small (1973). "The crystal structure of ammonium carbamate". Acta Crystallogr. B29 (11): 2317–2319. doi:10.1107/S056774087300662X.
  10. ^ Goldberg, R. N. Apparent Equilibrium Constants for Enzyme-catalyzed reactions (2009). CRC Handbook of Chemistry and Physics, 7–19. Retrieved from
  11. ^ Phosphorus Compounds: Advances in Research and Application: 2011 Edition
  12. ^ a b c 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. doi:10.1039/C0GC00674B.
  13. ^ United States Environmental Protection Agency. (2006). Inert Reassessment-Ammonium Carbamate [Data File]. Retrieved from
  14. ^ Mladen Litvić, Mirela Filipan, Ivan Pogorelić and Ivica Cepanec (2005): "Ammonium carbamate; mild, selective and efficient ammonia source for preparation of β-amino-α,β-unsaturated esters at room temperature". Green Chemistry, volume 7, issue 11, pages 771-774. doi:10.1039/B510276F