Ammonium

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Ammonium
Ammonium-2D.svg
Ammonium-3D-balls.png
Ammonium-3D-vdW.png
Names
IUPAC name
Ammonium ion
Systematic IUPAC name
Azanium [1]
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
MeSH D000644
PubChem CID
UNII
  • InChI=1S/H3N/h1H3/p+1
    Key: QGZKDVFQNNGYKY-UHFFFAOYSA-O
  • InChI=1/H3N/h1H3/p+1
    Key: QGZKDVFQNNGYKY-IKLDFBCSAZ
  • [NH4+]
Properties
[NH4]+
Molar mass 18.039 g·mol−1
Acidity (pKa)9.25
Conjugate base Ammonia
Structure
Tetrahedral
Related compounds
Other cations
Related compounds
Ammonium radical •NH4
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The ammonium cation is a positively charged polyatomic ion with the chemical formula NH+4 or [NH4]+. It is formed by the protonation of 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, the human impact in recent years could have an effect on the biological communities that depend on it.

Contents

Acid–base properties

Fumes from hydrochloric acid and ammonia forming a white cloud of ammonium chloride Hydrochloric acid ammonia.jpg
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 Bildung Ammonium.svg
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 ions 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−heat→NH3 + H2O

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

NH3(g) + HCl(aq) → [NH4]Cl(s)

Ammonia, when passed through CuSO4 (copper(II) sulfate) solution, changes its color from blue to deep blue, forming Schweizer's reagent.

CuSO4(aq) + 4 NH3(aq) + 4 H2O → [Cu(NH3)4(H2O)2](OH)2(aq) + H2SO4(aq)

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

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

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

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

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

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 = 175 pm)[ citation needed ] resembles the caesium cation (rionic = 183 pm).[ citation needed ]

Organic ions

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]+:

Dimethylammonium-formation-2D.png

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. Nitrogen Cycle 2.svg
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.

Because nitrogen often limits net primary production due to its use in enzymes that mediate the biochemical reactions that are necessary for life, ammonium is utilized by some microbes and plants. [5] For example, energy is released by the oxidation of ammonium in a process known as nitrification, which produces nitrate and nitrite. [6] This process is a form of autotrophy that is common amongst Nitrosomonas , Nitrobacter , Nitrosolobus , and Nitrosospira , amongst others. [6]

The amount of ammonium in soil that is available for nitrification by microbes varies depending on environmental conditions. [7] [8] For example, ammonium is deposited as 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. [9] Its availability in soils is also influenced by mineralization, which makes more ammonium available from organic nitrogen sources, and immobilization, which sequesters ammonium into organic nitrogen sources, both of which are mitigated by biological factors. [6]

Conversely, nitrate and nitrite can be reduced to ammonium as a way for living organisms to access nitrogen for growth in a process known as assimilatory nitrate reduction. [10] Once assimilated, it can be incorporated into proteins and DNA. [11]

Ammonium can accumulate in soils where nitrification is slow or inhibited, which is common in hypoxic soils. [12] For example, ammonium mobilization is one of the key factors for the symbiotic association between plants and fungi, called mycorrhizae. [13] However, plants that consistently utilize ammonium as a nitrogen source often must invest into more extensive root systems due to ammonium's limited mobility in soils compared to other nitrogen sources. [14] [15]

Human Impact

Ammonium deposition from the atmosphere has increased in recent years due to volatilization from livestock waste and increased fertilizer use. [16] 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. [17]

Metal

The ammonium cation has very similar properties to the heavier alkali metal cations and is often considered a close equivalent. [18] [19] [20] Ammonium is expected to behave as a metal ([NH4]+ ions in a sea of electrons) at very high pressures, such as inside giant planets such as Uranus and Neptune. [19] [20]

Under normal conditions, ammonium does not exist as a pure metal but does as an amalgam (alloy with mercury). [21]

See also

Related Research Articles

<span class="mw-page-title-main">Acid</span> Chemical compound giving a proton or accepting an electron pair

An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H+), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.

<span class="mw-page-title-main">Ammonia</span> Chemical compound

Ammonia is an inorganic chemical compound of nitrogen and hydrogen with the formula NH3. A stable binary hydride and the simplest pnictogen hydride, ammonia is a colourless gas with a distinctive pungent smell. Biologically, it is a common nitrogenous waste, and it contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to fertilisers. Around 70% of ammonia produced industrially is used to make fertilisers in various forms and composition, such as urea and diammonium phosphate. Ammonia in pure form is also applied directly into the soil.

<span class="mw-page-title-main">Amine</span> Chemical compounds and groups containing nitrogen with a lone pair (:N)

In chemistry, amines are compounds and functional groups that contain a basic nitrogen atom with a lone pair. Amines are formally derivatives of ammonia, wherein one or more hydrogen atoms have been replaced by a substituent such as an alkyl or aryl group. Important amines include amino acids, biogenic amines, trimethylamine, and aniline. Inorganic derivatives of ammonia are also called amines, such as monochloramine.

Urea, also called carbamide, is an organic compound with chemical formula CO(NH2)2. This amide has two amino groups joined by a carbonyl functional group. It is thus the simplest amide of carbamic acid.

<span class="mw-page-title-main">Nitrogen cycle</span> Biogeochemical cycle by which nitrogen is converted into various chemical forms

The nitrogen cycle is the biogeochemical cycle by which nitrogen is converted into multiple chemical forms as it circulates among atmospheric, terrestrial, and marine ecosystems. The conversion of nitrogen can be carried out through both biological and physical processes. Important processes in the nitrogen cycle include fixation, ammonification, nitrification, and denitrification. The majority of Earth's atmosphere (78%) is atmospheric nitrogen, making it the largest source of nitrogen. However, atmospheric nitrogen has limited availability for biological use, leading to a scarcity of usable nitrogen in many types of ecosystems.

<span class="mw-page-title-main">Base (chemistry)</span> Type of chemical substance

In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.

<span class="mw-page-title-main">Hydroxylamine</span> Inorganic compound

Hydroxylamine is an inorganic compound with the chemical formula NH2OH. The compound is in a form of a white hygroscopic crystals. Hydroxylamine is almost always provided and used as an aqueous solution. It is consumed almost exclusively to produce Nylon-6. The oxidation of NH3 to hydroxylamine is a step in biological nitrification.

<span class="mw-page-title-main">Nitrification</span> Biological oxidation of ammonia/ammonium to nitrate

Nitrification is the biological oxidation of ammonia to nitrate via the intermediary nitrite. Nitrification is an important step in the nitrogen cycle in soil. The process of complete nitrification may occur through separate organisms or entirely within one organism, as in comammox bacteria. The transformation of ammonia to nitrite is usually the rate limiting step of nitrification. Nitrification is an aerobic process performed by small groups of autotrophic bacteria and archaea.

Nitrogen assimilation is the formation of organic nitrogen compounds like amino acids from inorganic nitrogen compounds present in the environment. Organisms like plants, fungi and certain bacteria that can fix nitrogen gas (N2) depend on the ability to assimilate nitrate or ammonia for their needs. Other organisms, like animals, depend entirely on organic nitrogen from their food.

<span class="mw-page-title-main">Imidazole</span> Chemical compound

Imidazole (ImH) is an organic compound with the formula C3N2H4. It is a white or colourless solid that is soluble in water, producing a mildly alkaline solution. In chemistry, it is an aromatic heterocycle, classified as a diazole, and has non-adjacent nitrogen atoms in meta-substitution.

<span class="mw-page-title-main">Ammonium sulfate</span> Chemical compound

Ammonium sulfate (American English and international scientific usage; ammonium sulphate in British English); (NH4)2SO4, is an inorganic salt with a number of commercial uses. The most common use is as a soil fertilizer. It contains 21% nitrogen and 24% sulfur.

Acid salts are a class of salts that produce an acidic solution after being dissolved in a solvent. Its formation as a substance has a greater electrical conductivity than that of the pure solvent. An acidic solution formed by acid salt is made during partial neutralization of diprotic or polyprotic acids. A half-neutralization occurs due to the remaining of replaceable hydrogen atoms from the partial dissociation of weak acids that have not been reacted with hydroxide ions to create water molecules.

The Kjeldahl method or Kjeldahl digestion (Danish pronunciation:[ˈkʰelˌtɛˀl]) in analytical chemistry is a method for the quantitative determination of a sample's organic nitrogen plus ammonia/ammonium. (NH3/NH4+). Without modification, other forms of inorganic nitrogen, for instance nitrate, are not included in this measurement. Using an empirical relation between Kjeldahl nitrogen and protein, it is an important method for indirectly quantifying protein content of a sample. This method was developed by Johan Kjeldahl in 1883.

The chemical element nitrogen is one of the most abundant elements in the universe and can form many compounds. It can take several oxidation states; but the most common oxidation states are -3 and +3. Nitrogen can form nitride and nitrate ions. It also forms a part of nitric acid and nitrate salts. Nitrogen compounds also have an important role in organic chemistry, as nitrogen is part of proteins, amino acids and adenosine triphosphate.

In chemistry, an onium ion is a cation formally obtained by the protonation of mononuclear parent hydride of a pnictogen, chalcogen, or halogen. The oldest-known onium ion, and the namesake for the class, is ammonium, NH+4, the protonated derivative of ammonia, NH3.

<span class="mw-page-title-main">Hexaamminecobalt(III) chloride</span> Chemical compound

Hexaamminecobalt(III) chloride is the chemical compound with the formula [Co(NH3)6]Cl3. It is the chloride salt of the coordination complex [Co(NH3)6]3+, which is considered an archetypal "Werner complex", named after the pioneer of coordination chemistry, Alfred Werner. The cation itself is a metal ammine complex with six ammonia ligands attached to the cobalt(III) ion.

<span class="mw-page-title-main">Ammonium dinitramide</span> Chemical compound

Ammonium dinitramide (ADN) is an inorganic compound with the chemical formula [NH4][N(NO2)2]. It is the ammonium salt of dinitraminic acid HN(NO2)2. It consists of ammonium cations [NH4]+ and dinitramide anions −N(NO2)2. ADN decomposes under heat to leave only nitrogen, oxygen, and water.

In chemistry, a fatty amine is loosely defined as any amine possessing a mostly linear hydrocarbon chain of eight or more carbon atoms. They are typically prepared from the more abundant fatty acids, with vegetable or seed-oils being the ultimate starting material. As such they are often mixtures of chain lengths, ranging up to about C22. They can be classified as oleochemicals. Commercially important members include coco amine, oleylamine, tallow amine, and soya amine. These compounds and their derivatives are used as fabric softeners, froth flotation agents, corrosion inhibitors, lubricants and friction modifiers. They are also the basis for a variety of cosmetic formulations.

<span class="mw-page-title-main">Ammonia pollution</span> Chemical contamination

Ammonia pollution is pollution by the chemical ammonia (NH3) – a compound of nitrogen and hydrogen which is a byproduct of agriculture and industry. Common forms include air pollution by the ammonia gas emitted by rotting agricultural slurry and fertilizer factories while natural sources include the burning coal mines of Jharia, the caustic Lake Natron and the guano of seabird colonies. Gaseous ammonia reacts with other pollutants in the air to form fine particles of ammonium salts, which affect human breathing. Ammonia gas can also affect the chemistry of the soil on which it settles and will, for example, degrade the conditions required by the sphagnum moss and heathers of peatland.

<span class="mw-page-title-main">Nitrogen pentahydride</span> Chemical compound

Nitrogen pentahydride, also known as ammonium hydride is a hypothetical compound with the chemical formula NH5. There are two theoretical structures of nitrogen pentahydride. One structure is trigonal bipyramidal molecular geometry type NH5 molecule. Its nitrogen atom and hydrogen atoms are covalently bounded, and its symmetry group is D3h. Another predicted structure of nitrogen pentahydride is an ionic compound, composed of an ammonium ion and a hydride ion (NH4+H−). Until now, no one has synthesized this substance, or proved its existence, and related experiments have not directly observed nitrogen pentahydride. It is only speculated that it may be a reactive intermediate based on reaction products. Theoretical calculations show this molecule is thermodynamically unstable. The reason might be similar to the instability of nitrogen pentafluoride, so the possibility of its existence is low. However, nitrogen pentahydride might exist in special conditions or high pressure. Nitrogen pentahydride was considered for use as a solid rocket fuel for research in 1966.

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. ↑ Schlesinger, William H.; Bernhardt, Emily S. (2020-01-01), Schlesinger, William H.; Bernhardt, Emily S. (eds.), "Chapter 12 - The Global Cycles of Nitrogen, Phosphorus and Potassium", Biogeochemistry (Fourth Edition), 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.
  4. ↑ "Ammonium Salts". VIAS Encyclopedia.
  5. ↑ Schlesinger, William H.; Bernhardt, Emily S. (2020-01-01), Schlesinger, William H.; Bernhardt, Emily S. (eds.), "Chapter 12 - The Global Cycles of Nitrogen, Phosphorus and Potassium", Biogeochemistry (Fourth Edition), 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
  6. 1 2 3 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.
  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. ↑ Campbell, Neil A.; Reece, Jane B. (2002). Biology. Internet Archive. San Francisco : Benjamin Cummings. ISBN   978-0-8053-6624-2.
  10. ↑ Tiedje, J. M.; Sørensen, J.; Chang, Y.-Y. L. (1981). "Assimilatory and Dissimilatory Nitrate Reduction: Perspectives and Methodology for Simultaneous Measurement of Several Nitrogen Cycle Processes". Ecological Bulletins (33): 331–342. ISSN   0346-6868. JSTOR   45128674.
  11. ↑ 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.
  12. ↑ Wang, Lixin; Macko, Stephen A. (March 2011). "Constrained preferences in nitrogen uptake across plant species and environments". Plant, Cell & Environment. 34 (3): 525–534. doi:10.1111/j.1365-3040.2010.02260.x. ISSN   0140-7791. PMID   21118424.
  13. ↑ Hodge, Angela; Storer, Kate (2015-01-01). "Arbuscular mycorrhiza and nitrogen: implications for individual plants through to ecosystems". Plant and Soil. 386 (1): 1–19. Bibcode:2015PlSoi.386....1H. doi:10.1007/s11104-014-2162-1. ISSN   1573-5036.
  14. ↑ Raven, John A.; Linda, Bernd Wollenweber; Handley, L. (May 1992). "Ammonia and ammonium fluxes between photolithotrophs and the environment in relation to the global nitrogen cycle". New Phytologist. 121 (1): 5–18. doi:10.1111/j.1469-8137.1992.tb01087.x. ISSN   0028-646X.
  15. ↑ Bloom, A. J.; Jackson, L. E.; Smart, D. R. (March 1993). "Root growth as a function of ammonium and nitrate in the root zone". Plant, Cell & Environment. 16 (2): 199–206. doi:10.1111/j.1365-3040.1993.tb00861.x. ISSN   0140-7791.
  16. ↑ Ackerman, Daniel; Millet, Dylan B.; Chen, Xin (January 2019). "Global Estimates of Inorganic Nitrogen Deposition Across Four Decades". Global Biogeochemical Cycles. 33 (1): 100–107. Bibcode:2019GBioC..33..100A. doi:10.1029/2018GB005990. ISSN   0886-6236.
  17. ↑ 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. doi: 10.3389/fpls.2022.894365 . ISSN   1664-462X. PMC   9251499 . PMID   35795351.
  18. ↑ 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
  19. 1 2 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.
  20. 1 2 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 .
  21. ↑ 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.
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