Wikipedia

Ammonium carbamate

Ammonium carbamate is a chemical compound with the formula [NH4][H2NCO2] consisting of ammonium cation NH+4 and carbamate anion NH2COO−. 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 NH3 with carbon dioxide CO2, and will slowly decompose to those gases at ordinary temperatures and pressures. It is an intermediate in the industrial synthesis of urea (NH2)2CO, an important fertilizer.[4]

Ammonium carbamate
Names
IUPAC name
Ammonium carbamate
Other names
hartshorn, sal volatile, ammonium amidocarbonate, ammonium aminoformate, [1]
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.012.896 Edit this at Wikidata
EC Number
  • 214-185-2
14637 (G)
RTECS number
  • EY8575000
UNII
UN number 9083
  • InChI=1S/CH3NO2.H3N/c2-1(3)4;/h2H2,(H,3,4);1H3
  • [O-]C(=O)N.[NH4+]
Properties
[NH4]NH2CO2
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)
Thermochemistry
-642.5 kJ/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Harmful if ingested, harmful to aquatic life, harmful if inhaled, respiatory tract irritation, skin irritation, eye irritation
GHS labelling:
GHS07: Exclamation mark
Warning
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
0
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
Safety data sheet (SDS) External MSDS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Properties edit

Solid-gas equilibrium edit

In a closed container solid ammonium carbamate is in equilibrium with carbon dioxide and ammonia [5][6][7]

[NH4][NH2CO2] ⇌ 2 NH3 + CO2

Lower temperatures shift the equilibrium towards the carbamate.

At higher temperatures ammonium carbamate condenses into urea:

[NH4][NH2CO2] → (NH2)2CO + H2O

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.[6]

Equilibrium in water edit

At ordinary temperatures and pressures, ammonium carbamate exists in aqueous solutions as an equilibrium with ammonia and carbon dioxide, and the anions bicarbonate, HCO−3, and carbonate, CO2−3.[8][6][9] Indeed, solutions of ammonium carbonate or bicarbonate will contain some carbamate anions too.

H2NCO−2 + 2H2O ⇌ NH+4 + HCO−3 + OH−
H2NCO−2 + H2O ⇌ NH+4 + CO2−3

Structure edit

The structure of solid ammonium carbamate has been confirmed by X-ray crystallography. The oxygen centers form hydrogen bonds to the ammonium cation.[10] There are two polymorphs, α and β, both in the orthorhombic crystal system but differing in their space group. The α polymorph is in space group Pbca (no. 61), whereas the β polymorph is in Ibam (no. 72). The α polymorph is more volatile.[11]

Natural occurrence edit

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:[12][13]

ATP + [NH2CO2][NH4] → ADP + H2NC(O)OPO2−3

Preparation edit

From liquid ammonia and dry ice edit

Ammonium carbamate is prepared by the direct reaction between liquid ammonia and dry ice (solid carbon dioxide):[5]

2 NH3 + CO2 → [NH2CO2][NH4]

From gaseous ammonia and carbon dioxide edit

Ammonium carbamate can be prepared by reaction of the two gases at high temperature (175–225 °C) and high pressure (150–250 bar).[14]

It can also be obtained by bubbling gaseous CO2 and NH3 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.[14]

Uses edit

Urea synthesis edit

Ammonium carbamate is an intermediate in the industrial production of urea. A typical industrial plant that makes urea can produce up to 4000 tons a day.[15] in this reactor and can then be dehydrated to urea according to the following equation:[14]

[NH2CO2][NH4] → (NH2)2CO + H2O

Pesticide formulations edit

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 phosphine formed by a hydrolysis reaction.[16]

Laboratory edit

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.[17]

Preparation of metal carbamates edit

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 NH2COO−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]

References edit

  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. ^ Jäger, Peter; Rentzea, Costin N.; Kieczka, Heinz (2000). "Carbamates and Carbamoyl Chlorides". ULLMANN'S Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a05_051. ISBN 3527306730.
  5. ^ a b Brooks, L. A.; Audrieta, L. F.; Bluestone, H.; Jofinsox, W. C. (1946). "Ammonium Carbamate". Inorganic Syntheses. Vol. 2. pp. 85–86. doi:10.1002/9780470132333.ch23. ISBN 9780470132333.
  6. ^ a b c K. G. Clark; V. L. Gaddy; C. E. Rist (October 1933). "Equilibria in the Ammonium Carbamate-Urea-Water System". Industrial and engineering chemistry. 25 (10): 1092–1096. doi:10.1021/IE50286A008. ISSN 0019-7866. Wikidata Q59410838.
  7. ^ 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
  8. ^ 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
  9. ^ Fabrizio Mani, Maurizio Peruzzini, and Piero Stoppioni (2006): "CO2 absorption by aqueous NH3 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
  10. ^ J. M. Adams; R. W. H. Small (15 November 1973). "The crystal structure of ammonium carbamate". Acta crystallographica. Section B: Structural crystallography and crystal chemistry. 29 (11): 2317–2319. doi:10.1107/S056774087300662X. ISSN 0567-7408. Wikidata Q59410837.
  11. ^ Kuhn, Norbert; Ströbele, Markus; Meyer, H.-Jürgen (2007). "Über die Identität eines sogenannten Ammoniumcarbonat-Präparates". Z. Anorg. Allg. Chem. 633 (4): 635–656. doi:10.1002/zaac.200600392.
  12. ^ 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 Archived 2016-03-04 at the Wayback Machine
  13. ^ Phosphorus Compounds: Advances in Research and Application: 2011 Edition. ScholarlyEditions. January 9, 2012. ISBN 9781464925573 – via Google Books.
  14. ^ a b c Francesco Barzagli; Fabrizio Mani; Maurizio Peruzzini (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. ISSN 1463-9262. Wikidata Q59410840.
  15. ^ "Dangote fertiliser plant and other fertilizer plants operating in Nigeria". 22 March 2022.
  16. ^ United States Environmental Protection Agency. (2006). Inert Reassessment-Ammonium Carbamate [Data File]. Retrieved from http://www.epa.gov/opprd001/inerts/carbamate.pdf
  17. ^ 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
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