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This article is about the ammonium ion. For its neutral conjugate base, see ammonia.
This article is about the molecular ion. For the ancient city, see Siwa Oasis.

Ammonium is a modified form of ammonia that has an extra hydrogen atom. It is a positively charged (cationic) molecular ion with the chemical formula NH+4 or [NH4]+. It is formed by the addition of a proton (a hydrogen nucleus) to ammonia (NH3). Ammonium is also a general name for positively charged (protonated) substituted amines and quaternary ammonium cations ([NR4]+), where one or more hydrogen atoms are replaced by organic or other groups (indicated by R). Not only is ammonium a source of nitrogen and a key metabolite for many living organisms, but it is an integral part of the global nitrogen cycle.[2] As such, human impact in recent years could have an effect on the biological communities that depend on it.

Acid–base properties

Fumes from hydrochloric acid and ammonia forming a white cloud of ammonium chloride

The ammonium ion is generated when ammonia, a weak base, reacts with Brønsted acids (proton donors):

H+ + NH3 → [NH4]+

The ammonium ion is mildly acidic, reacting with Brønsted bases to return to the uncharged ammonia molecule:

[NH4]+ + B− → HB + NH3

Thus, the treatment of concentrated solutions of ammonium salts with a strong base gives ammonia. When ammonia is dissolved in water, a tiny amount of it converts to ammonium ions:

H2O + NH3 ⇌ OH− + [NH4]+

The degree to which ammonia forms the ammonium ion depends on the pH of the solution. If the pH is low, the equilibrium shifts to the right: more ammonia molecules are converted into ammonium ions. If the pH is high (the concentration of hydrogen ions is low and hydroxide ions is high), the equilibrium shifts to the left: the hydroxide ion abstracts a proton from the ammonium ion, generating ammonia.

Formation of ammonium compounds can also occur in the vapor phase; for example, when ammonia vapor comes in contact with hydrogen chloride vapor, a white cloud of ammonium chloride forms, which eventually settles out as a solid in a thin white layer on surfaces.

Salts and characteristic reactions

Formation of ammonium

Ammonium cation is found in a variety of salts such as ammonium carbonate, ammonium chloride, and ammonium nitrate. Most simple ammonium salts are very soluble in water. An exception is ammonium hexachloroplatinate, the formation of which was once used as a test for ammonium. The ammonium salts of nitrate and especially perchlorate are highly explosive, in these cases, ammonium is the reducing agent.

In an unusual process, ammonium radicals (NH4) form an amalgam. Such species are prepared by the addition of sodium amalgam to a solution of ammonium chloride.[3] This amalgam eventually decomposes to release ammonia and hydrogen.[4]

To find whether the ammonium ion is present in the salt, first, the salt is heated in presence of alkali hydroxide releasing a gas with a characteristic smell, which is ammonia.

[NH4]+ + OH− → NH3 + H2O

To further confirm ammonia, it is passed through a glass rod dipped in an HCl solution (hydrochloric acid), creating white dense fumes of ammonium chloride.

NH3 + HCl → [NH4]Cl

Ammonia or ammonium ion when added to Nessler’s reagent gives a brown color precipitate known as the iodide of Million’s base in basic medium.

Ammonium ion when added to chloroplatinic acid gives a yellow precipitate of ammonium hexachloroplatinate(IV).

H2[PtCl6] + [NH4]+ → [NH4]2[PtCl6](s) + 2 H+

Ammonium ion when added to sodium cobaltinitrite gives a yellow precipitate of ammonium cobaltinitrite.

Na3[Co(NO2)6] + 3 [NH4]+ → [NH4]3[Co(NO2)6](s) + 3 Na+

Ammonium ion gives a white precipitate of ammonium bitartrate when added to potassium bitartrate.

KC4H5O6 + [NH4]+ → [NH4]C4H5O6(s) + K+

Structure and bonding

The lone electron pair on the nitrogen atom (N) in ammonia, represented as a line above the N, forms a coordinate bond with a proton (H+). After that, all four N−H bonds are equivalent, being polar covalent bonds. The ion has a tetrahedral structure and is isoelectronic with methane and the borohydride anion. In terms of size, the ammonium cation (rionic = 148 pm), it is intermediate in size between potassium (144 pm) and rubidium ions (152 pm) when in an octahedral environment.[5]

The vibrational spectrum consists of two main sets of absorptions, νN-H and δHNH. These bands are found near 3300 and 1400 cm−1.[6]

Organic ions

See also: Amine

The hydrogen atoms in the ammonium ion can be substituted with an alkyl group or some other organic group to form a substituted ammonium ion (IUPAC nomenclature: aminium ion). Depending on the number of organic groups, the ammonium cation is called a primary, secondary, tertiary, or quaternary. Except the quaternary ammonium cations, the organic ammonium cations are weak acids.

An example of a reaction forming an ammonium ion is that between dimethylamine, (CH3)2NH, and an acid to give the dimethylammonium cation, [(CH3)2NH2]+:

Quaternary ammonium cations have four organic groups attached to the nitrogen atom, they lack a hydrogen atom bonded to the nitrogen atom. These cations, such as the tetra-n-butylammonium cation, are sometimes used to replace sodium or potassium ions to increase the solubility of the associated anion in organic solvents. Primary, secondary, and tertiary ammonium salts serve the same function but are less lipophilic. They are also used as phase-transfer catalysts and surfactants.

An unusual class of organic ammonium salts is derivatives of amine radical cations, [•NR3]+ such as tris(4-bromophenyl)ammoniumyl hexachloroantimonate.

Biology

Ammonium exists as a result of ammonification and decomposers. Ammonium is eventually nitrified, where it contributes to the flow of nitrogen through the ecosystem. Human impacts are not shown here, but can impact the global nitrogen cycle.

Ammonium is utilized by living organism for biosynthesis and as an electron donor.[2]

In part related to the use of fertilizers, there is much interest in ammonium in soils.[7][8]

Ammonia oxidation

oxidation of ammonium occurs nitrification, which produces nitrate and nitrite.[9] This process is a form of autotrophy that is common amongst Nitrosomonas, Nitrobacter, Nitrosolobus, and Nitrosospira, amongst others.[9] Conversely, nitrate and nitrite can be reduced to ammonium in a process known as assimilatory nitrate reduction.[10]

When nitrification is slow, which is common in hypoxic soils, ammonium can accumulate in soils.[11]

Excretion

Ammonium is a waste product from some animals, although it is converted into urea in mammals, sharks, and amphibians, and into uric acid in birds, reptiles, and terrestrial snails.[12] If transport or excretion of ammonium is interrupted, accumulation of ammonium can lead to hyperammonemia.[13] Toxic effects arise from high concentrations of glutamate, which can interfere with the TCA cycle.[14]

Biosynthesis with ammonium

Once assimilated, it can be incorporated into proteins and DNA.[15]

Ammonium transporters

transportors or carriers occur widely.[16] Common human transport proteins for ammonium include the Amt/Mep/Rh family.[17] [18][19] Glutamine synthetase is one participant.[20][21]

Urea Cycle showing the formation of carbamoyl phosphate in the mitochondria and the rest of the Urea Cycle occurring in the cytoplasm of the liver cell.

Other mechanisms include the Cahill cycle, glutamate, and glutamine to transport ammonium into the mitochondria.[22] This ammonium combines with bicarbonate to form carbamoyl phosphate which then enters the urea cycle to be excreted as urea.[20]

Human impact

Ammonium deposition from the atmosphere has increased in recent years due to volatilization from livestock waste and increased fertilizer use.[23] Because net primary production is often limited by nitrogen, increased ammonium levels could impact the biological communities that rely on it. For example, increasing nitrogen content has been shown to increase plant growth, but aggravate soil phosphorus levels, which can impact microbial communities.[24]

Metal

The ammonium cation has very similar properties to the heavier alkali metal cations and is often considered a close equivalent.[25][26][27] Neutral ammonium is expected to behave as a metal ([NH4]+ ions in a sea of electrons) at very high pressures, such as inside the giant planets Uranus and Neptune.[26][27]

Under normal conditions, ammonium does not exist as a pure metal but does as an amalgam (NH4 alloy with mercury).[28]

See also

References

  1. ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. pp. 71,105,314. Electronic version.
  2. ^ a b Schlesinger, William H.; Bernhardt, Emily S. (2020-01-01). “Chapter 12 – The Global Cycles of Nitrogen, Phosphorus and Potassium”. In Schlesinger, William H.; Bernhardt, Emily S. (eds.). Biogeochemistry (Fourth ed.). Academic Press. pp. 483–508. doi:10.1016/b978-0-12-814608-8.00012-8. ISBN 978-0-12-814608-8. Retrieved 2024-03-08.
  3. ^ “Pseudo-binary compounds”. Archived from the original on 2020-07-27. Retrieved 2007-10-12.
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  7. ^ Barsdate, Robert J.; Alexander, Vera (January 1975). “The Nitrogen Balance of Arctic Tundra: Pathways, Rates, and Environmental Implications”. Journal of Environmental Quality. 4 (1): 111–117. Bibcode:1975JEnvQ…4..111B. doi:10.2134/jeq1975.00472425000400010025x. ISSN 0047-2425.
  8. ^ Nadelhoffer, Knute J.; Aber, John D.; Melillo, Jerry M. (1984-10-01). “Seasonal patterns of ammonium and nitrate uptake in nine temperate forest ecosystems”. Plant and Soil. 80 (3): 321–335. Bibcode:1984PlSoi..80..321N. doi:10.1007/BF02140039. ISSN 1573-5036.
  9. ^ a b Rosswall, T. (1982). “Microbiological regulation of the biogeochemical nitrogen cycle / Regulación microbiana del ciclo bíogeoquímico del nitrógeno”. Plant and Soil. 67 (1/3): 15–34. doi:10.1007/BF02182752. ISSN 0032-079X. JSTOR 42934020.
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  14. ^ Visek, Willard J. (March 1984). “Ammonia: Its Effects on Biological Systems, Metabolic Hormones, and Reproduction”. Journal of Dairy Science. 67 (3): 481–498. doi:10.3168/jds.S0022-0302(84)81331-4. ISSN 0022-0302. PMID 6371080.
  15. ^ Llácer, José L; Fita, Ignacio; Rubio, Vicente (2008-12-01). “Arginine and nitrogen storage”. Current Opinion in Structural Biology. Catalysis and regulation / Proteins. 18 (6): 673–681. doi:10.1016/j.sbi.2008.11.002. hdl:10261/111022. ISSN 0959-440X. PMID 19013524.
  16. ^ Mohiuddin, Shamim S.; Khattar, Divya (2025), “Biochemistry, Ammonia”, StatPearls, Treasure Island (FL): StatPearls Publishing, PMID 31082083, retrieved 2026-02-22
  17. ^ Williamson, Gordon; Bizior, Adriana; Harris, Thomas; Pritchard, Leighton; Hoskisson, Paul A.; Javelle, Arnaud (2024-01-31). “Biological ammonium transporters from the Amt/Mep/Rh superfamily: mechanism, energetics, and technical limitations”. Bioscience Reports. 44 (1) BSR20211209. doi:10.1042/BSR20211209. ISSN 0144-8463. PMC 10794816. PMID 38131184.
  18. ^ Ariz, Idoia; Boeckstaens, Mélanie; Gouveia, Catarina; Martins, Ana Paula; Sanz-Luque, Emanuel; Fernández, Emilio; Soveral, Graça; von Wirén, Nicolaus; Marini, Anna M.; Aparicio-Tejo, Pedro M.; Cruz, Cristina (2018-09-07). “Nitrogen isotope signature evidences ammonium deprotonation as a common transport mechanism for the AMT-Mep-Rh protein superfamily”. Science Advances. 4 (9) eaar3599. Bibcode:2018SciA….4.3599A. doi:10.1126/sciadv.aar3599. ISSN 2375-2548. PMC 6135547. PMID 30214933.
  19. ^ Weiner, I. David; Hamm, L. Lee (2007-03-01). “Molecular Mechanisms of Renal Ammonia Transport”. Annual Review of Physiology. 69 (1): 317–340. doi:10.1146/annurev.physiol.69.040705.142215. ISSN 0066-4278. PMC 4313553. PMID 17002591.
  20. ^ a b Adeva, Maria M.; Souto, Gema; Blanco, Natalia; Donapetry, Cristóbal (November 2012). “Ammonium metabolism in humans”. Metabolism. 61 (11): 1495–1511. doi:10.1016/j.metabol.2012.07.007. PMID 22921946.
  21. ^ Jakhar, Deepika; Sarin, Shiv K.; Kaur, Savneet (2024-12-04). “Gut microbiota and dynamics of ammonia metabolism in liver disease”. npj Gut and Liver. 1 (1): 11. doi:10.1038/s44355-024-00011-x. ISSN 3004-9806.
  22. ^ PubChem. “Glucose-Alanine Cycle”. pubchem.ncbi.nlm.nih.gov. Retrieved 2026-04-09.
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  24. ^ Dong, Junfu; Cui, Xiaoyong; Niu, Haishan; Zhang, Jing; Zhu, Chuanlu; Li, Linfeng; Pang, Zhe; Wang, Shiping (2022-06-20). “Effects of Nitrogen Addition on Plant Properties and Microbiomes Under High Phosphorus Addition Level in the Alpine Steppe”. Frontiers in Plant Science. 13 894365. doi:10.3389/fpls.2022.894365. ISSN 1664-462X. PMC 9251499. PMID 35795351.
  25. ^ Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
  26. ^ a b Stevenson, D. J. (November 20, 1975). “Does metallic ammonium exist?”. Nature. 258 (5532): 222–223. Bibcode:1975Natur.258..222S. doi:10.1038/258222a0. S2CID 4199721.
  27. ^ a b Bernal, M. J. M.; Massey, H. S. W. (February 3, 1954). “Metallic Ammonium”. Monthly Notices of the Royal Astronomical Society. 114 (2): 172–179. Bibcode:1954MNRAS.114..172B. doi:10.1093/mnras/114.2.172.
  28. ^ Reedy, J.H. (October 1, 1929). “Lecture demonstration of ammonium amalgam”. Journal of Chemical Education. 6 (10): 1767. Bibcode:1929JChEd…6.1767R. doi:10.1021/ed006p1767.