{{otheruses}}
{{chembox new
| ImageFileL1 = Ammonia-2D-dimensions.png
| ImageSizeL1 = 150 px
| ImageFileR1 = Ammonia-3D-vdW.png
| ImageSizeR1 = 150 px
| IUPACName = Azane NH3
| OtherNames = Ammonia <br/>Hydrogen nitride<br/>Spirit of Hartshorn<br />Nitro-Sil<br />Vaporole <ref>[http://webbook.nist.gov/cgi/cbook.cgi?Name=Ammonia Ammonia data at NIST WebBook], last accessed May 7 2007.</ref>
| Section1 = {{Chembox Identifiers
| CASNo = 7664-41-7
| PubChem = 222
| SMILES = N
| InChI = 1/H3N/h1H3
| RTECS = BO0875000
}}
| Section2 = {{Chembox Properties
| Formula = NH₃
| MolarMass = 17.0306 g/mol
| Appearance = Colorless gas with strong pungent odor
| Density = 0.6942 <ref>[http://webbook.nist.gov/cgi/fluid.cgi?Action=Load&ID=C7664417&Type=IsoTherm&PLow=0.9&PHigh=1.1&PInc=0.1&T=25&RefState=DEF&TUnit=C&PUnit=bar&DUnit=kg%2Fm3&HUnit=kJ%2Fmol&WUnit=m%2Fs&VisUnit=uPa*s&STUnit=N%2Fm NIST Chemistry WebBook] (website page of the National Institute of Standards and Technology) URL last accessed May 15 2007</ref>
| MeltingPt = -77.73 °C (195.42 [[Kelvin|K]])
| Melting_notes =
| BoilingPt = -33.34 °C (239.81 K)
| Boiling_notes =
| Solubility = 89.9 g/100 [[Milliliter|mL]] at 0 °[[Celsius|C]]
| SolubleOther =
| Solvent =
| pKa =
| pKa =
| pKb = 4.75 (reaction with H₂O)
| RefractIndex = [[Dielectric constant|ε_r]]
}}
| Section3 = {{Chembox Structure
| MolShape = Terminus
| Dipole = 1.42 [[Debye|D]]
}}
| Section7 = {{Chembox Hazards
| MainHazards = Hazardous gas, caustic, corrosive
| NFPA-H = 3
| NFPA-F = 1
| NFPA-R =
| NFPA-O =
| RPhrases = {{R10}}, {{R23}}, {{R34}}, {{R50}}<br/> {{S1/2}}, {{S16}}, {{S36/37/39}},<br />{{S45}}, {{S61}}
| SPhrases =
| RSPhrases =
| FlashPt = None<ref>[http://www.wdserviceco.com/03aug06MSDS/msdsANH.pdf MSDS Sheet] from W.D. Service Co.</ref>
| Autoignition = 651 °C
| ExploLimits =
| PEL = }}
| Section8 = {{Chembox Related
| OtherAnions = [[ammonium hydroxide|hydroxide]] (NH₄OH)
| OtherCations = [[Ammonium]] (NH₄⁺)
| OtherFunctn = [[ammonium chloride|chloride]] (NH₄Cl)
| Function =
| OtherCpds = [[Hydrazine]]<br />[[Hydrazoic acid]]<br />[[Hydroxylamine]]<br />[[Chloramine]] }}
}}
'''Ammonia''' is a [[chemical compound|compound]] with the [[chemical formula|formula]] [[nitrogen|N]][[hydrogen|H₃]]. It is normally encountered as a [[gas]] with a characteristic pungent [[odor]]. Ammonia contributes significantly to the nutritional needs of the planet as a precursor to foodstuffs and fertilizers. Ammonia, either directly or indirectly, is a building block for the synthesis of many pharmaceuticals. Although in wide use, ammonia is caustic and hazardous.

Ammonia, used commercially is usually named ''anhydrous ammonia.'' This term emphasizes the absence of water. Because NH₃ boils at -33 °C, the liquid must be stored under pressure or at low temperature. Its [[heat of vaporization]] is, however, sufficiently high that NH₃ can be readily handled in ordinary beakers in a [[fume hood]]. "Household ammonia" or "[[ammonium hydroxide]]" is a solution of NH₃ in water. The strength of such solutions is measured in units of [[baume]] ([[density]]), with 26 degrees baume (about 30 weight percent ammonia at 15.5 °C) being the typical high concentration commercial product.<ref name=LaRoche>[http://www.airgasspecialtyproducts.com/UserFiles/laroche/PDF/AAPhysical.pdf Ammonium hydroxide physical properties]</ref> Household ammonia ranges in concentration from 5 to 10 weight percent ammonia. See [[Baumé scale]].

==Structure and basic chemical properties==
The ammonia molecule has a [[Trigonal pyramid (chemistry)|trigonal pyramid]] shape, as predicted by [[VSEPR theory]]. The [[nitrogen]] atom in the molecule has a [[Lone pair|lone electron pair]], and ammonia acts as a [[Base (chemistry)|base]], a proton acceptor. This shape gives the molecule an overall [[dipole]] moment and makes it [[Polar molecule|polar]] so that ammonia readily dissolves in [[Water (molecule)|water]]. In water, a very small percentage of NH₃ is converted into the [[ammonium]] [[cation]] (NH₄⁺). Thus, the term ammonium hydroxide is a misnomer. The degree to which ammonia forms the ammonium ion increases upon lowering the [[pH]] of the [[solution]]&mdash; at "physiological" pH (~7), about 99% of the ammonia molecules are [[Protonation|protonated]]. Temperature and salinity also affect the proportion of NH₄⁺. NH₄⁺ has the shape of a regular [[tetrahedron]].

The main use of ammonia is for [[fertilizer]] (83% in 2003). Another major application is its conversion to [[explosive]]s, because nitric acid is made via oxidation of ammonia. The ''entire'' nitrogen content of all manufactured [[organic compound]]s is derived from ammonia.<ref name=Appl>Max Appl “Ammonia” in Ullmann's Encyclopedia of Industrial Chemistry Wiley-VCH Verlag; Weinheim, 2002.DOI: 10.1002/14356007.a02_143.pub2</ref>

===Natural occurence===
Ammonia is found in small quantities in the atmosphere, being produced from the [[putrefaction]] of nitrogenous animal and vegetable matter. Ammonia and ammonium salts are also found in small quantities in rainwater, whereas [[ammonium chloride]] (sal-ammoniac), and [[ammonium sulfate]] are found in volcanic districts; crystals of [[ammonium bicarbonate]] have been found in [[Patagonia]]n [[guano]]. The kidneys secrete NH₃ to neutralize excess acid.<ref name= >[http://www.ncbi.nlm.nih.gov/entrez/query.fcgi?cmd=Retrieve&db=PubMed&list_uids=10360635&dopt=Urine electrolytes and the urine anion and osmolar gaps.]</ref> Ammonium salts also are found distributed through all fertile soil and in seawater. Substances containing ammonia, or those that are similar to it, are called ''ammoniacal''.

==History==
The Romans called the [[ammonium chloride]] deposits they collected from near the Temple of Jupiter [[Amun]] ([[Greek language|Greek]] Ἄμμων ''Ammon'') in [[ancient Libya]] 'sal ammoniacus' (salt of Amun) because of proximity to the nearby temple<ref>{{cite web | title = Ammonia | work = h2g2 Eponyms | publisher = BBB.CO.UK | date = January 11, 2003 | url = http://www.bbc.co.uk/dna/h2g2/alabaster/A632990| accessdate = 2007-11-08 }}</ref>. Salts of ammonia have been known from very early times; thus the term ''Hammoniacus sal''<ref name="Mineral Data">[http://webmineral.com/data/Sal-ammoniac.shtml Webmineral website] URL last accessed August 27 2006</ref> appears in the writings of [[Pliny the Elder|Pliny]], although it is not known whether the term is identical with the more modern ''sal-ammoniac''.<ref name="Mineral Data"/>

In the form of sal-ammoniac, ammonia was known to the [[alchemy|alchemists]] as early as the 13th century, being mentioned by [[Albertus Magnus]].<ref name="astronomy">[http://www.absoluteastronomy.com/ref/ammonia Absolouteastronomy.com] URL last accessed April 24 2006</ref> It was also used by [[Dye|dyer]]s in the [[Middle Ages]] in the form of fermented [[urine]]<ref name="astronomy"/> to alter the colour of vegetable dyes. In the 15th century, [[Basilius Valentinus]] showed that ammonia could be obtained by the action of alkalis on sal-ammoniac. At a later period, when sal-ammoniac was obtained by distilling the hoofs and horns of oxen and neutralizing the resulting carbonate with [[hydrochloric acid]], the name "spirit of hartshorn" was applied to ammonia.<ref name="astronomy"/>

Gaseous ammonia was first isolated by [[Joseph Priestley]] in 1774 and was termed by him ''alkaline air''; however it was acquired by the alchemist [[Basil Valentine]].<ref name="Alchemy" > Abraham, Lyndy. Marvell and alchemy. Aldershot Scolar 1990.</ref> Eleven years later in 1785, [[Claude Louis Berthollet]] ascertained its composition.

The [[Haber process]] to produce ammonia from the nitrogen in the air was developed by [[Fritz Haber]] and [[Carl Bosch]] in 1909 and patented in 1910. It was first used on an industrial scale by the Germans during [[World War I]],<ref name=Appl> following the allied blockade that cut off the supply of nitrates from [[Chile]]. The ammonia was used to produce explosives to sustain their war effort.<ref name="Conquering" >Smith, Roland. Conquering Chemistry 2001</ref>

==Synthesis and production==
Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Dozens of [[chemical plant]]s worldwide produce ammonia. The worldwide [[ammonia production]] in 2004 was 109 million [[metric tonnes]].<ref name="usgs">[http://minerals.usgs.gov/minerals/pubs/commodity/nitrogen/nitromcs05.pdf United States Geological Survey publication]</ref> The [[People's Republic of China]] produced 28.4% of the worldwide production followed by [[India]] with 8.6%, [[Russia]] with 8.4%, and the [[United States]] with 8.2%.<ref name="usgs"/> About 80% or more of the ammonia produced is used for fertilizing agricultural crops.<ref name="usgs"/>

Before the start of [[World War I]], most ammonia was obtained by the dry [[distillation]]<ref>[http://nobelprize.org/chemistry/laureates/1918/press.html Nobel Prize in Chemistry (1918)] - [[Haber process]]. URL last accessed April 24 2006</ref> of nitrogenous vegetable and animal waste products, including [[camel]] [[manure|dung]], where it was [[distillation|distilled]]<ref name="Conquering"/> by the reduction of [[nitrous acid]] and [[nitrite]]s with [[hydrogen]]; in addition, it was produced by the distillation of [[coal]],<ref name="Conquering"/> and also by the decomposition of ammonium salts by [[alkaline]] hydroxides<ref>[http://www.bbc.co.uk/dna/h2g2/A1002934 BBC.co.uk] URL last accessed April 24 2006</ref> such as [[calcium oxide|quicklime]], the salt most generally used being the chloride ([[ammonium chloride|sal-ammoniac]]) thus:

::2 NH₄Cl + 2 CaO → CaCl₂ + Ca(OH)₂ + 2 NH₃

Today, the typical modern ammonia-producing plant first converts [[natural gas]] (i.e., [[methane]]) or [[liquified petroleum gas]] (such gases are [[propane]] and [[butane]]) or petroleum [[naphtha]] into gaseous [[hydrogen]]. Starting with a natural gas feedstock, the processes used in producing the hydrogen are:

* The first step in the process entails removal of [[sulfur]] compounds from the feedstock, because sulfur deactivates the [[catalyst]]s used in subsequent steps. Catalytic [[hydrogenation]] converts organosulfur compounds into gaseous [[hydrogen sulfide]]:
::H₂ + RSH → RH + H₂S(''g'')

*The hydrogen sulfide is then removed by passing the gas through beds of [[zinc oxide]] where it is absorbed and converted to solid [[zinc sulfide]]:
::H₂S + ZnO → ZnS + H₂O

* Catalytic [[steam reforming]] of the sulfur-free feedstock is then used to form hydrogen plus [[carbon monoxide]]:
::CH₄ + H₂O → CO + 3 H₂

* In the next step, the [[water gas shift reaction]] is used to convert the [[carbon monoxide]] into [[carbon dioxide]] and more hydrogen:
::CO + H₂O → CO₂ + H₂

* The carbon dioxide is then removed either by absorption in aqueous [[ethanolamine]] solutions or by [[adsorption]] in [[Pressure Swing Adsorption|pressure swing adsorbers]] (PSA) using proprietary solid adsorption media.

* The final step in producing the hydrogen is to use catalytic methanation to remove any small residual amounts of carbon monoxide or carbon dioxide from the hydrogen:
::CO + 3 H₂ → CH₄ + H₂O
::CO₂ + 4 H₂ → CH₄ + 2 H₂O

* To produce the desired end-product ammonia, the hydrogen is then catalytically reacted with nitrogen (derived from process air) to form anhydrous liquid ammonia. This step is known as the ammonia synthesis loop (also referred to as the [[Haber-Bosch]] process):

::3 H₂ + N₂ → 2 NH₃

The steam reforming, shift conversion, carbon dioxide removal and methanation steps each operate at absolute pressures of about 25 to 35 [[Bar (unit)|bar]], and the ammonia synthesis loop operates at absolute pressures ranging from 60 to 180 bar, depending upon which proprietary design is used. There are many engineering and construction companies that offer proprietary designs for ammonia synthesis plants. [[Haldor Topsoe]] of [[Denmark]], [[Lurgi AG]] of [[Germany]], [[ThyssenKrupp|Uhde]] of [[Germany]], and [[Kellogg, Brown and Root]] of the United States are among the most experienced companies in that field.<ref>[http://www.highbeam.com/doc/1G1:54711794/Grassroots+success+with+KAAP~R~(Kellogg+Brown+and+Roots+Advanced+Ammonia+Process).html?refid=SEO Kellogg Brown's Ammonia Process] URL last accessed April 24 2006</ref>

As the availability and usage of fossil fuel become problematic (see [[peak oil]] and [[climate change]]), the [[hydrogen]] required for ammonia synthesis could in principle be obtained from [[electrolysis]] (currently 4% of hydrogen production is from electrolysis) or thermal chemical cracking of [[water]], but these alternatives are presently impractical. The heat needed for thermal cracking can be obtained from nuclear reaction, while the electricity needed for electrolysis can be obtained from various renewable energy sources such as [[Wind turbine|wind]], [[Photovoltaics|solar]], [[hydroelectricity]], and various forms of [[ocean energy]] especially that of [[OTEC]]. On really windy days on wind farms the power lines can't handle all the electricity. A possible use for the excess electricity would be to use electrolysis on water to acquire the needed hydrogen. Again alternatives to the production of ammonia from natural gas and air are uneconomic and the environmental benefits have not been established.

==Biosynthesis==
In certain organisms, ammonia is produced from atmospheric N₂ by [[enzyme]]s called [[nitrogenase]]s. The overall process is called [[nitrogen fixation]]. Although it is unlikely that biomimetic methods will be developed that are competitive with the [[Haber process]], intense effort has been directed toward understanding the mechanism of biological nitrogen fixation. The scientific interest in this problem is motivated by the unusual structure of the active site of the enzyme, which consists of an Fe₇MoS₉ ensemble.

Ammonia is also a metabolic product of [[amino acid]] [[deamination]]. In humans, it is quickly converted to [[urea]], which is much less toxic. This urea is a major component of the dry weight of [[urine]].

== Properties ==
Ammonia is a colorless [[gas]] with a characteristic pungent smell similar to [[human urine]], as urine decomposes to release ammonia. It is [[lighter than air]], its density being 0.589 times that of [[Earth's atmosphere|air]]. It is easily liquefied due to the strong hydrogen bonding between molecules; the [[liquid]] boils at -33.3 °C, and solidifies at -77.7 °C to a mass of white crystals. [[Liquid]] ammonia possesses strong [[ion]]izing powers ([[Dielectric constant|ε]] = 22), and [[solution]]s of [[salt]]s in liquid ammonia have been much studied. Liquid ammonia has a very high [[standard enthalpy change of vaporization]] (23.35&nbsp;[[Joule|kJ]]/mol, ''cf.'' [[water (molecule)|water]] 40.65&nbsp;kJ/mol, [[methane]] 8.19&nbsp;kJ/mol, [[phosphine]] 14.6&nbsp;kJ/mol) and can therefore be used in laboratories in non-insulated vessels at room temperature, even though it is well above its boiling point.

It is [[miscible]] with water. Ammonia in an aqueous solution can be expelled by boiling. The [[water|aqueous]] solution of ammonia is [[Base (chemistry)|basic]]. The maximum concentration of ammonia in water (a [[saturation (chemistry)|saturated]] solution) has a [[density]] of 0.880 g /[[cubic centimetre|cm³ ]] and is often known as '.880 Ammonia'. Ammonia does not burn readily or sustain [[combustion]], except under narrow fuel to air mixtures from 15-25% air. When mixed with [[oxygen]], it burns with a pale yellowish-green flame. At high temperature and in the presence of a suitable catalyst, ammonia is decomposed into its constituent elements. [[Chlorine]] catches fire when passed into ammonia, forming nitrogen and [[hydrochloric acid]]; unless the ammonia is present in excess, the highly explosive [[nitrogen trichloride]] (NCl₃) is also formed.

The ammonia molecule readily undergoes [[nitrogen inversion]] at room temperature - that is, the nitrogen atom passes through the [[plane of symmetry]] of the three hydrogen atoms; a useful analogy is an [[umbrella]] turning itself inside out in a strong wind. The energy barrier to this inversion is 24.7 kJ/mol in ammonia, and the [[resonance frequency]] is 23.79 [[Hertz|GHz]], corresponding to [[microwave]] radiation of a [[wavelength]] of 1.260 cm. The absorption at this frequency was the first [[Microwave spectroscopy|microwave spectrum]] to be observed.<ref name="Cleeton">C. E. Cleeton & N. H. Williams, 1934 - [http://prola.aps.org/abstract/PR/v45/i4/p234_1 Online version; archive]. URL last accessed May 8, 2006</ref>

=== Formation of salts ===
One of the most characteristic properties of ammonia is its power of combining directly with [[acid]]s to form [[salt]]s; thus with [[hydrochloric acid]] it forms [[ammonium chloride]] (sal-ammoniac); with [[nitric acid]], [[ammonium nitrate]], etc. However perfectly dry ammonia will not combine with perfectly dry [[hydrogen chloride]], a gas, moisture being necessary to bring about the reaction.<ref>Baker, H. B. (1894). ''J. Chem. Soc.'' '''65''': 612.</ref>
::NH₃ + [[Hydrochloric acid|HCl]] → [[Ammonium chloride|NH₄Cl]]

The salts produced by the action of ammonia on acids are known as the [[:Category:Ammonium compounds|ammonium salts]] and all contain the [[ammonium]] [[ion]] (NH₄⁺).

=== Acidity ===
Although ammonia is well-known as a base, it can also act as an extremely weak [[acid]]. It is a protic substance, and is capable of dissociation into the '''amide''' (NH₂^&minus;) ion, for example when solid [[lithium nitride]] is added to liquid ammonia, forming a [[lithium amide]] solution:

::Li₃N_(''s'')+ 2 NH_{3 (''l'')} → 3 Li⁺_(''am'') + 3 NH₂^&minus;_(''am'')

This is a [[Brønsted-Lowry]] acid-base reaction in which ammonia is acting as an acid.

=== Formation of other compounds ===
In [[organic chemistry]], ammonia can act as a [[nucleophile]] in [[Nucleophilic substitution|substitution]] reactions. [[Amine]]s can be formed by the reaction of ammonia with [[alkyl halide]]s, although the resulting &ndash;NH₂ group is also nucleophilic and secondary and tertiary amines are often formed as by-products. An excess of ammonia helps minimise multiple substitution, and neutralises the [[hydrogen halide]] formed. [[Methylamine]] is prepared commercially by the reaction of ammonia with [[chloromethane]], and the reaction of ammonia with 2-bromopropanoic acid has been used to prepare [[racemic]] [[alanine]] in 70% yield. [[Ethanolamine]] is prepared by a ring-opening reaction with [[ethylene oxide]]: the reaction is sometimes allowed to go further to produce [[diethanolamine]] and [[triethanolamine]].

[[Amide]]s can be prepared by the reaction of ammonia with a number of [[carboxylic acid]] derivatives. [[Acyl chloride]]s are the most reactive, but the ammonia must be present in at least a twofold excess to neutralise the [[hydrogen chloride]] formed. [[Ester]]s and [[anhydride]]s also react with ammonia to form amides. Ammonium salts of carboxylic acids can be [[Dehydration|dehydrated]] to amides so long as there are no thermally sensitive groups present: temperatures of 150&ndash;200 °C are required.

The [[hydrogen]] in ammonia is capable of replacement by [[metal]]s, thus [[magnesium]] burns in the gas with the formation of [[magnesium nitride]] Mg₃N₂, and when the gas is passed over heated [[sodium]] or [[potassium]], sodamide, NaNH₂, and potassamide, KNH₂, are formed. Where necessary in [[IUPAC nomenclature|substitutive nomenclature]], [[IUPAC]] recommendations prefer the name '''azane''' to ammonia: hence [[chloramine]] would be named ''chloroazane'' in substitutive nomenclature, not ''chloroammonia''.

=== Ammonia as a ligand ===
[[Image:TetraAmmineDiAquaCopper.png‎|thumb|right|200px|[[Ball-and-stick model]] of the tetraamminediaquacopper(II) cation, [Cu(NH₃)₄(H₂O)₂]²⁺]]

[[Image:Diamminesilver(I)-3D-balls.png|thumb|right|200px|Ball-and-stick model of the diamminesilver(I) cation, [Ag(NH₃)₂]⁺]]

Ammonia can act as a [[ligand]] in [[transition metal]] [[complex (chemistry)|complexes]]. It is a pure σ-donor, in the middle of the [[spectrochemical series]], and shows intermediate [[HSAB concept|hard-soft]] behaviour. For historical reasons, ammonia is named '''ammine''' in the nomenclature of [[coordination compound]]s. Some notable ammine complexes include:
*'''Tetraamminediaquacopper(II)''', [Cu(NH₃)₄(H₂O)₂]²⁺, a characteristic dark blue complex formed by adding ammonia to solution of copper(II) salts.
*'''Diamminesilver(I)''', [Ag(NH₃)₂]⁺, the active species in [[Tollens' reagent]]. Formation of this complex can also help to distinguish between precipitates of the different silver halides: [[Silver chloride|AgCl]] is soluble in dilute (2M) ammonia solution, [[Silver bromide|AgBr]] is only soluble in concentrated ammonia solution while [[Silver iodide|AgI]] is insoluble in aqueous solution of ammonia.

Ammine complexes of [[chromium]](III) were known in the late 19th century, and formed the basis of [[Alfred Werner]]'s theory of coordination compounds. Werner noted that only two isomers (''fac''- and ''mer''-) of the complex [CrCl₃(NH₃)₃] could be formed, and concluded that the ligands must be arranged around the metal ion at the [[wikt:vertex|vertices]] of an [[octahedron]]. This has since been confirmed by [[X-ray crystallography]].

An ammine ligand bound to a metal ion is markedly more [[acid]]ic than a free ammonia molecule, although deprotonation in aqueous solution is still rare. One example is the [[Mercury(I) chloride|Calomel reaction]], where the resulting amidomercury(II) compound is highly insoluble.
::Hg₂Cl₂ + 2 NH₃ → Hg + HgCl(NH₂) + NH₄⁺ + Cl^&minus;

==Uses==
===Nitric Acid production===
The most important single use of ammonia is in the production of [[nitric acid]]. A mixture of one part ammonia to nine parts air is passed over a [[platinum]] gauze [[catalyst]] at 850 °C, whereupon the ammonia is [[oxidisation|oxidized]] to [[nitric oxide]].

::4 NH₃ + 5 O₂ → 4 NO + 6 H₂O
::2 NO + O₂ → 2 NO₂
::2 NO₂ + 2 H₂O → 2 HNO₃ + H₂

The catalyst is essential, as the normal oxidation (or combustion) of ammonia gives [[Nitrogen|dinitrogen]] and water: the production of nitric oxide is an example of [[kinetic control]]. As the gas mixture cools to 200&ndash;250 °C, the nitric oxide is in turn oxidized by the excess of [[oxygen]] present in the mixture, to give [[nitrogen dioxide]]. This is reacted with water to give nitric acid for use in the production of [[fertilizer]]s and [[explosive]]s.

===Fertilizer===
In addition to serving as a fertilizer ingredient, ammonia can also be used directly as a fertilizer by forming a solution with irrigation water, without additional chemical processing. This later use allows the continuous growing of nitrogen dependent crops such as [[maize]] (corn) without [[agriculture|crop]] rotation but this type of use leads to poor [[soil]] health.

===Refrigeration===
Ammonia's thermodynamic properties made it one of the [[refrigerants]] commonly used in [[refrigeration]] units prior to the discovery of [[dichlorodifluoromethane]]<ref>[http://www.wou.edu/~avorder/Refrigeration.htm Dichlorodifluoromethane] by Aaron Vorderstrasse, Western Oregon University.</ref> in 1928, also known as [[Freon]] or R12.

But ammonia is toxic, gaseous, irritant, and corrosive to [[copper]] [[alloy]]s, and over a kilo is needed for even a miniature fridge. With an ammonia refrigerant, the ever present risk of an escape brings with it a risk to life. However data on ammonia escapes has shown this to be an extremely small risk in practice, and there is consequently no control on the use of ammonia refrigeration in densely populated areas and buildings in almost all jurisdictions in the world.

Its use in domestic [[refrigeration]] has been mostly replaced by CFCs and HFCs in the first world, which are more or less non-toxic and non-[[flammable]], and butane and propane in the 3rd world, which despite their high [[flammable|flammability]] do not seem to have produced any significant level of accidents. Ammonia has continued to be used for miniature and multifuel fridges, such as minibars and caravan refridgerators.

These ammonia absorption cycle domestic refrigerators do not use compression and expansion cycles, but are driven by temperature differences. However the [[energy efficiency]] of such refrigerators is relatively low. Today the smallest refrigerators mostly use solid state peltier [[thermopile]] heat pumps rather than the ammonia absorption cycle.

Ammonia continues to be used as a [[refrigerant]] in large industrial processes such as bulk icemaking and industrial food processing. Since the implication of [[haloalkane]]s being major contributors to [[ozone depletion]], ammonia is again seeing increasing use as a refrigerant. Ammonia is increasingly popular in commercial applications, such as in grocery store freezer cases and refrigerated displays.

===Disinfectant===
It is also sometimes added to drinking water along with [[chlorine]] to form [[chloramine]], a [[disinfectant]]. Unlike chlorine alone, chloramine does not combine with organic (carbon containing) materials to form [[carcinogen]]ic [[halomethane]]s such as [[chloroform]]. However, chlorine and ammonia should never be mixed in an uncontrolled environment because they cause a chemical reaction that releases toxic gas. See [[Ammonia#Safety precautions|Safety precautions]] for more information.

===Fuel===
Ammonia was used during [[World War II]] fuel shortages to power buses in Belgium and used in engine and solar energy applications prior to 1900. Liquid ammonia was used as the fuel of the rocket airplane, the [[X-15]]. Although not as powerful as other fuels, it left no soot in the reusable rocket engine and its density approximately matches that for the oxidizer, liquid oxygen, which simplified the aircraft's design. Ammonia is proposed as a practical, clean ([[Carbon dioxide|CO₂]]-free), alternative to [[fossil fuel]] for [[internal combustion engines]].<ref>http://www.energy.iastate.edu/Renewable/ammonia/ammonia/ammoniaMtg07.htm</ref> In 1981 a Canadian company converted a 1981 Chevrolet Impala to operate using ammonia as fuel.<ref>http://www.youtube.com/watch?v=L0hBAz6MxC4</ref><ref>http://www.gregvezina.com</ref>
Ammonia is marketed as a low-emission fuel.<ref>http://nh3fuel.com</ref>

===Cigarettes===
During the 1960s, [[tobacco]] companies such as [[Brown & Williamson]] and [[Philip Morris]] began using ammonia in [[cigarette]]s.{{Fact|date=December 2007}} Many hypotheses have appeared in the media and technical literature that ammonia enhances the amount of nicotine available to the smoker, nicotine's bioavailability, and the reinforcing or addictive ability<ref>
Alix M. Freedman, "[http://www.pulitzer.org/year/1996/national-reporting/works/impact.html 'Impact Booster': Tobacco Firm Shows How Ammonia Spurs Delivery of Nicotine]", ''[[The Wall Street Journal]]'', Dec. 28, 1995.</ref>. In contrast, a number of recent publications in the scientific literature have demonstrated that ammonia-forming compounds in tobacco and ammonia in mainstream tobacco smoke do not increase either the total amount or total rate of nicotine to the bloodstream or brains of smokers.<ref>Jeffrey Seeman, "[http://pubs.acs.org/cgi-bin/abstract.cgi/crtoec/2007/20/i03/abs/tx600290v.html Possible Role of Ammonia on the Deposition, Retention, and Absorption of Nicotine in Humans while Smoking]", Chem. Res. Toxicol., 20 (3), 326 -343, 2007. 10.1021/tx600290v S0893-228x(60)00290-7, 2007.</ref>.

===Illicit Drug Manufacture===
* Largely before the popularization of [[Crack_cocaine | crack]] [[cocaine]], [[Ammonium hydroxide]] was commonly used in the production of "[[Cocaine#Freebase | freebase cocaine]]"
* Conversely (when noting the presence or absence of water in the solution), '''Anhydrous ammonia''' is often used in one of the most dangerous methods for the [[Methamphetamine#Illicit_production | production of methamphetamine]]
== Ammonia's role in biologic systems and human disease ==
Ammonia is an important source of nitrogen for living systems. Although atmospheric nitrogen abounds, few living creatures are capable of utilizing this nitrogen. Nitrogen is required for the synthesis of amino acids, which are the building blocks of [[protein]]. Some plants rely on ammonia and other nitrogenous wastes incorporated into the soil by decaying matter. Others, such as nitrogen-fixing [[legume]]s, benefit from [[symbiosis|symbiotic]] relationships with [[rhizobia]] which create ammonia from atmospheric nitrogen.<ref>M.B. Adjei, K.H. Quesenberry and C.G. Chamblis. ''Nitrogen Fixation and Inoculation of Forage Legumes'' [http://edis.ifas.ufl.edu/AG152 University of Florida IFAS Extension] June 2002.</ref>

Ammonia also plays a role in both normal and abnormal animal [[physiology]]. Ammonia is created through normal amino acid metabolism and is toxic in high concentrations.<ref>[http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?sid=10323059 PubChem Substance Summary], last accessed May 7 2007</ref> The [[liver]] converts ammonia to [[urea]] through a series of reactions known as the [[urea cycle]]. Liver dysfunction, such as that seen in [[cirrhosis]], may lead to elevated amounts of ammonia in the blood ([[hyperammonemia]]). Likewise, defects in the enzymes responsible for the urea cycle, such as [[ornithine transcarbamylase]], lead to hyperammonemia. Hyperammonemia contributes to the confusion and [[coma]] of [[hepatic encephalopathy]] as well as the neurologic disease common in people with urea cycle defects and [[organic aciduria]]s.<ref>Zschocke, Johannes, and Georg Hoffman. ''Vademecum Metabolism.'' Friedrichsdorf, Germany: Milupa GmbH, 2004.</ref>

Ammonia is important for normal animal acid/base balance. After formation of ammonium from [[glutamine]], [[α-ketoglutarate]] may be degraded to produce two molecules of [[bicarbonate]] which are then available as buffers for dietary acids. Ammonium is excreted in the urine resulting in net acid loss. Ammonia may itself diffuse across the renal tubules, combine with a hydrogen ion, and thus allow for further acid excretion.<ref>Rose, Burton, and Helmut Rennke. ''Renal Pathophysiology.'' Baltimore, Maryland: Williams & Wilkins, 1994.</ref>

===Theoretical role in alternative biochemistry===
Ammonia has been proposed as a possible replacement for water as a bodily solvent in the theoretical [[alternative biochemistries]] of lifeforms that do not use [[carbon]] for cellular structure and [[water]] as a solvent to dissolve bodily solutes and allow essential parts of metabolic processes to occur. It has been suggested that ammonia would be most favorable for lifeforms that live in temperatures below the freezing point of water{{Fact|date=December 2007}}.

== Liquid ammonia as a solvent ==
:''See also: [[Inorganic nonaqueous solvent]]''
Liquid ammonia is the best-known and most widely studied non-aqueous ionizing solvent. Its most conspicuous property is its ability to dissolve alkali metals to form highly coloured, electrically conducting solutions containing [[solvated electron]]s. Apart from these remarkable solutions, much of the chemistry in liquid ammonia can be classified by analogy with related reactions in aqueous solutions. Comparison of the physical properties of NH₃ with those of water shows that NH₃ has the lower melting point, boiling point, density, [[viscosity]], [[dielectric constant]] and [[electrical conductivity]]; this is due at least in part to the weaker H bonding in NH₃ and the fact that such bonding cannot form cross-linked networks since each NH₃ molecule has only 1 lone-pair of electrons compared with 2 for each H₂O molecule. The ionic self-[[dissociation constant]] of liquid NH₃ at &minus;50 °C is approx. 10^-33 mol²·l^-2.

=== Solubility of salts ===
{|
|-
! &nbsp;
! Solubility (g of salt per 100 g liquid NH₃)
|-
| [[Ammonium acetate]]
| 253.2
|-
| [[Ammonium nitrate]]
| 389.6
|-
| [[Lithium nitrate]]
| 243.7
|-
| [[Sodium nitrate]]
| 97.6
|-
| [[Potassium nitrate]]
| 10.4
|-
| [[Sodium fluoride]]
| 0.35
|-
| [[Sodium chloride]]
| 3.0
|-
| [[Sodium bromide]]
| 138.0
|-
| [[Sodium iodide]]
| 161.9
|-
| [[Sodium thiocyanate]]
| 205.5
|-
|}

Liquid ammonia is an ionizing solvent, although less so than water, and dissolves a range of ionic compounds including many [[nitrate]]s, [[nitrite]]s, [[cyanide]]s and [[thiocyanate]]s. Most [[ammonium]] salts are soluble, and these salts act as [[acid]]s in liquid ammonia solutions. The solubility of [[halide]] salts increases from [[fluoride]] to [[iodide]]. A saturated solution of [[ammonium nitrate]] contains 0.83 mol solute per mole of ammonia, and has a [[vapour pressure]] of less than 1&nbsp;bar even at 25 °C.

=== Solutions of metals ===
:''See also: [[Solvated electron]], [[metallic solution]]''
Liquid ammonia will dissolve the [[alkali metal]]s and other [[Electronegativity|electropositive]] metals such as [[calcium]], [[strontium]], [[barium]], [[europium]] and [[ytterbium]]. At low concentrations (<0.06 mol/L), deep blue solutions are formed: these contain metal cations and [[solvated electron]]s, free electrons which are surrounded by a cage of ammonia molecules.

These solutions are very useful as strong reducing agents. At higher concentrations, the solutions are metallic in appearance and in electrical conductivity. At low temperatures, the two types of solution can coexist as [[Wiktionary:immiscible|immiscible]] phases.

=== Redox properties of liquid ammonia ===
:''See also: [[Redox]].''
{| cellpadding="5"
|-
! &nbsp;
! [[Standard electrode potential|''E''°]] (V, ammonia)
! [[Standard electrode potential|''E''°]] (V, water)
|-
| Li⁺ + e^&minus; {{unicode|&#8652;}} Li
|style="text-align: center;"| &minus;2.24
|style="text-align: center;"| &minus;3.04
|-
| K⁺ + e^&minus; {{unicode|&#8652;}} K
|style="text-align: center;"| &minus;1.98
|style="text-align: center;"| &minus;2.93
|-
| Na⁺ + e^&minus; {{unicode|&#8652;}} Na
|style="text-align: center;"| &minus;1.85
|style="text-align: center;"| &minus;2.71
|-
| Zn²⁺ + 2e^&minus; {{unicode|&#8652;}} Zn
|style="text-align: center;"| &minus;0.53
|style="text-align: center;"| &minus;0.76
|-
| NH₄⁺ + e^&minus; {{unicode|&#8652;}} ½ H₂ + NH₃
|style="text-align: center;"| 0.00
|style="text-align: center;"| &ndash;
|-
| Cu²⁺ + 2e^&minus; {{unicode|&#8652;}} Cu
|style="text-align: center;"| +0.43
|style="text-align: center;"| +0.34
|-
| Ag⁺ + e^&minus; {{unicode|&#8652;}} Ag
|style="text-align: center;"| +0.83
|style="text-align: center;"| +0.80
|-
|}

The range of thermodynamic stability of liquid ammonia solutions is very narrow, as the potential for oxidation to [[Nitrogen|dinitrogen]], [[Standard electrode potential|''E''°]] (N₂ + 6NH₄⁺ + 6e^&minus; {{unicode|&#8652;}} 8NH₃), is only +0.04&nbsp;V. In practice, both oxidation to dinitrogen and reduction to [[Hydrogen|dihydrogen]] are slow. This is particularly true of reducing solutions: the solutions of the alkali metals mentioned above are stable for several days, slowly decomposing to the [[Amide|metal amide]] and dihydrogen. Most studies involving liquid ammonia solutions are done in reducing conditions: although oxidation of liquid ammonia is usually slow, there is still a risk of explosion, particularly if transition metal ions are present as possible catalysts.

== Detection and determination ==
Ammonia and ammonium salts can be readily detected, in very minute traces, by the addition of [[Nessler's solution]], which gives a distinct yellow coloration in the presence of the least trace of ammonia or ammonium salts. [[Sulfur sticks]] are burnt to detect small leaks in industrial ammonia refrigeration systems. Larger quantities can be detected by warming the salts with a caustic alkali or with [[calcium oxide|quicklime]], when the characteristic smell of ammonia will be at once apparent. The amount of ammonia in ammonium salts can be estimated quantitatively by distillation of the salts with [[sodium hydroxide|sodium]] or [[potassium hydroxide]], the ammonia evolved being absorbed in a known volume of standard [[sulfuric acid]] and the excess of acid then determined [[volumetric analysis|volumetrically]]; or the ammonia may be absorbed in [[hydrochloric acid]] and the [[ammonium chloride]] so formed precipitated as [[ammonium hexachloroplatinate]], (NH₄)₂PtCl₆.

===Interstellar space===
Ammonia was first detected in interstellar space in 1968, based on [[microwave]] emissions from the direction of the [[Milky Way|galactic core]].<ref> A.C. Cheung, D.M. Rank, C.H. Townes, D.D. Thornton, and W.J. Welch, 1968, "[http://adsabs.harvard.edu/cgi-bin/nph-bib_query?bibcode=1968PhRvL..21.1701C Detection of NH3 molecules in the interstellar medium by their microwave emission]," ''Phys. Rev. Lett.'' 21, 1701.</ref> This was the first [[polyatomic]] molecule to be so detected.
The sensitivity of the molecule to a broad range of excitations and the ease with which it can be observed in a number of regions has made ammonia one of the most important molecules for studies of [[molecular cloud]]s.<ref>P. T. P. Ho and C.H. Townes, 1983,
"[http://adsabs.harvard.edu/abs/1983ARA&A..21..239H Interstellar ammonia], ''Ann. Rev. Astron. Astrophys.'', vol. 21, pp. 239-70.</ref> The relative intensity of the ammonia lines can be used to measure the temperature of the emitting medium.

The following isotopic species of ammonia have been detected:
:NH₃, [[Isotopes of nitrogen|¹⁵N]]H₃, NH₂[[Deuterium|D]], NHD₂, and ND₃
The detection of triply-[[deuterium|deuterated]] ammonia was considered a surprise as deuterium is relatively scarce. It is thought that the low-temperature conditions allow this molecule to survive and accumulate.<ref>T. J. Millar, "Deuterium Fractionation in Interstellar Clouds", ''Space Science Reviews'', Vol. 106, Issue 1, pp 73-86.</ref> The ammonia molecule has also been detected in the atmospheres of the [[gas giant]] planets, including [[Jupiter (planet)|Jupiter]], along with other gases like [[methane]], [[hydrogen]], and [[helium]]. The interior of Saturn may include frozen crystals of ammonia.<ref>Edited by Kirk Munsell. Image page credit Lunar and Planetary Institute. NASA. "[http://solarsystem.nasa.gov/multimedia/display.cfm?IM_ID=166 NASA's Solar Exploration: Multimedia: Gallery: Gas Giant Interiors]". URL accessed April 26, 2006.</ref>
== Safety precautions ==
=== Toxicity and storage information ===
[[Image:Hydrochloric acid ammonia.jpg|thumb|Hydrochloric acid sample releasing HCl fumes which are reacting with ammonia fumes to produce a white smoke of ammonium chloride.]]
The toxicity of ammonia solutions does not usually cause problems for humans and other mammals, as a specific mechanism exists to prevent its build-up in the bloodstream. Ammonia is converted to [[carbamoyl phosphate]] by the enzyme [[carbamoyl phosphate synthase]], and then enters the [[urea cycle]] to be either incorporated into [[amino acid]]s or excreted in the urine. However [[fish]] and [[amphibian]]s lack this mechanism, as they can usually eliminate ammonia from their bodies by direct excretion. Ammonia even at dilute concentrations is highly toxic to aquatic animals, and for this reason it is [[Directive 67/548/EEC|classified]] as ''dangerous for the environment''. Ammonium compounds should never be allowed to come in contact with bases (unless an intended and contained reaction), as dangerous quantities of ammonia gas could be released.

=== Household use ===
Solutions of ammonia (5&ndash;10% by weight) are used as household cleaners, particularly for glass. These solutions are irritating to the eyes and [[mucous membrane]]s (respiratory and digestive tracts), and to a lesser extent the skin. They should '''never''' be mixed with [[chlorine]]-containing products or strong oxidants, for example household [[bleach]], as a variety of toxic and [[carcinogen]]ic compounds are formed (''e.g.'', [[chloramine]], [[hydrazine]], and chlorine gas). Ammonia and [[sodium hypochlorite]] react to form a number of products, depending on the temperature, concentration, and how they are mixed. <ref>{{Cite journal | last =Rizk-Ouaini | first =Rosette | author-link = | last2 =Ferriol, Michel; Gazet, Josette; Saugier-Cohen Adad, Marie Therese | first2 = | author2-link = | title = Oxidation reaction of ammonia with sodium hypochlorite. Production and degradation reactions of chloramines. | journal = Bulletin de la Societe Chimique de France | volume =4 | issue = | pages =512–21 | date =1986 | year = | url = | doi = | id = }}</ref>. The main reaction is chlorination of ammonia, first giving [[chloramine]] (NH₂Cl), then NHCl₂ and finally [[nitrogen trichloride]] (NCl₃). These materials are very irritating to eyes and lungs and are toxic above certain concentrations.

=== Laboratory use of ammonia solutions ===
The hazards of ammonia solutions depend on the concentration: "dilute" ammonia solutions are usually 5&ndash;10% by weight (<5.62 mol/L); "concentrated" solutions are usually prepared at >25% by weight. A 25% (by weight) solution has a density of 0.907 g/cm³, and a solution which has a lower density will be more concentrated. The [[Directive 67/548/EEC|European Union classification]] of ammonia solutions is given in the table.

<!-- EU Index no. 007-001-01-2 -->
{| class="wikitable"
|-
! [[Concentration]]<br />by weight
! Molarity
! Mass/Volume
! Classification
! [[List of R-phrases|R-Phrases]]
|-
| 5&ndash;10%
| 2.87&ndash;5.62 mol/L
| 48.9&ndash;95.7 g/L
| Irritant ('''Xi''')
| {{R36/37/38}}
|-
| 10&ndash;25%
| 5.62&ndash;13.29 mol/L
| 95.7&ndash;226.3 g/L
| Corrosive ('''C''')
| {{R34}}
|-
| >25%
| >13.29 mol/L
| >226.3 g/L
| Corrosive ('''C''')<br />Dangerous for<br />the environment ('''N''')
| {{R34}}, {{R50}}
|-
|}

:''[[List of S-phrases|S-Phrases]]: {{S1/2}}, {{S16}}, {{S36/37/39}}, {{S45}}, {{S61}}.''

The ammonia vapour from concentrated ammonia solutions is severely irritating to the eyes and the respiratory tract, and these solutions should only be handled in a fume hood. Saturated ("0.880") solutions can develop a significant pressure inside a closed bottle in warm weather, and the bottle should be opened with care: this is not usually a problem for 25% ("0.900") solutions.

Ammonia solutions should not be mixed with [[halogen]]s, as toxic and/or explosive products are formed. Prolonged contact of ammonia solutions with [[silver (element)|silver]], [[mercury (element)|mercury]] or [[iodide]] salts can also lead to explosive products: such mixtures are often formed in [[qualitative analysis|qualitative chemical analysis]], and should be acidified and diluted before disposal once the test is completed.

=== Laboratory use of anhydrous ammonia (gas or liquid) ===
<!-- EU index no. 007-001-00-5 -->
<!-- R10, 23, 34, 50; S1/2, 9, 16, 26, 36/37/39, 45, 61 -->
Anhydrous ammonia is classified as '''toxic''' ('''T''') and '''dangerous for the environment''' ('''N'''). The gas is flammable ([[autoignition temperature]]: 651 °C) and can form explosive mixtures with air (16&ndash;25%). The [[permissible exposure limit]] (PEL) in the United States is 50&nbsp;[[Parts per million|ppm]] (35&nbsp;mg/m³), while the [[IDLH]] concentration is estimated at 300&nbsp;ppm. Repeated exposure to ammonia lowers the sensitivity to the smell of the gas: normally the odour is detectable at concentrations of less than 0.5&nbsp;ppm, but desensitized individuals may not detect it even at concentrations of 100&nbsp;ppm. Anhydrous ammonia corrodes [[copper (element)|copper]]- and [[zinc (element)|zinc]]-containing [[alloy]]s, and so [[brass]] fittings should not be used for handling the gas. Liquid ammonia can also attack rubber and certain plastics.

Ammonia reacts violently with the halogens, and causes the explosive [[polymerization]] of [[ethylene oxide]]. It also forms explosive compounds with compounds of [[gold (element)|gold]], [[silver]], [[Mercury (element)|mercury]], [[germanium]] or [[tellurium]], and with [[stibine]]. Violent reactions have also been reported with [[acetaldehyde]], [[hypochlorite]] solutions, [[potassium ferricyanide]] and [[peroxide]]s.

==Safety==
The U. S. [[Occupational Safety and Health Administration|Occupational Safety and Health Administration (OSHA)]] has set a 15-minute exposure limit for gaseous ammonia of 35 ppm by volume in the environmental air and an 8-hour exposure limit of 25 ppm by volume.<ref name=ToxFaqSheet>[http://www.atsdr.cdc.gov/tfacts126.pdf Toxic FAQ Sheet for Ammonia] published by the [[Agency for Toxic Substances and Disease Registry]] (ATSDR), September 2004</ref> Exposure to very high concentrations of gaseous ammonia can result in lung damage and death.<ref name=ToxFaqSheet/> Although ammonia is regulated in the United States as a non-flammable gas, it still meets the definition of a material that is toxic by inhalation and requires a hazardous safety permit when transported in quantities greater than 13,248 L (3,500 gallons).<ref>[http://www.fmcsa.dot.gov/safety-security/hazmat/hm-permitting.htm#one Hazardous Materials (HM) Safety Permits] from the website of the [[United States Department of Transportation]] (DOT)</ref>

== See also ==
* [[Ammonia (data page)]]
* [[Ammonia production]]
* [[Chlorination]]
* [[Water purification]]

== References ==
<div class="references-small"><references/></div>

== Bibliography ==
*{{1911}}
* {{cite book | last=Greenwood |first=N. N. |coauthors = Earnshaw, A. | title=Chemistry of the Elements | edition = 2nd Edn. | location =Oxford | publisher=Butterworth-Heinemann | year=1997 | id=ISBN 0-7506-3365-4 }}
* {{cite book | last=Housecroft |first=C. E. |coauthors = Sharpe, A. | title=Inorganic Chemistry | location =Harlow (UK) | publisher=Prentice Education | year=2001 | id=ISBN 0-582-31080-6 }}
* {{cite book | editor =Bretherick, L. | title=Hazards in the Chemical Laboratory | edition = 4th Edn. | location =London | publisher=Royal Society of Chemistry | year=1986 | id=ISBN 0-85186-489-9 }}
*{{RubberBible53rd}}

== External links ==
* [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc04/icsc0414.htm International Chemical Safety Card 0414] (anhydrous ammonia)
* [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc02/icsc0215.htm International Chemical Safety Card 0215] (aqueous solutions)
* [http://www.npi.gov.au/database/substance-info/profiles/8.html National Pollutant Inventory - Ammonia]
* {{PubChemLink|222}}
* {{inFrench|url=www.inrs.fr|name=Institut national de recherche et de securite}}
* [http://www.ammoniaspills.org Emergency Response to Ammonia Fertilizer Releases (Spills)] for the Minnesota Department of Agriculture

*[http://www.cdc.gov/niosh/topics/ammonia National Institute for Occupational Safety and Health - Ammonia Page]

[[Category:Hydrogen compounds]]
[[Category:Hydrides]]
[[Category:Nitrogen compounds]]
[[Category:Bases]]
[[Category:Nitrogen metabolism]]
[[Category:Household chemicals]]
[[Category:Refrigerants]]
[[Category:Occupational safety and health]]

{{Link FA|hr}}

[[ar:أمونياك]]
[[bs:Amonijak]]
[[bg:Амоняк]]
[[ca:Amoníac]]
[[cs:Amoniak]]
[[cy:Amonia]]
[[da:Ammoniak]]
[[de:Ammoniak]]
[[et:Ammoniaak]]
[[el:Αμμωνία]]
[[es:Amoníaco]]
[[eo:Amoniako]]
[[eu:Amoniako]]
[[fa:آمونیاک]]
[[fr:Ammoniac]]
[[hr:Amonijak]]
[[io:Amoniako]]
[[id:Amonia]]
[[it:Ammoniaca]]
[[he:אמוניה]]
[[ku:Ammonîak]]
[[lv:Amonjaks]]
[[lt:Amoniakas]]
[[hu:Ammónia]]
[[mk:Амонијак]]
[[ms:Ammonia]]
[[nl:Ammoniak]]
[[ja:アンモニア]]
[[no:Ammoniakk]]
[[nn:Ammoniakk]]
[[pl:Amoniak]]
[[pt:Amoníaco]]
[[ro:Amoniac]]
[[ru:Аммиак]]
[[scn:Ammunìaca]]
[[simple:Ammonia]]
[[sk:Amoniak]]
[[sl:Amonijak]]
[[sr:Амонијак]]
[[sh:Amonijak]]
[[fi:Ammoniakki]]
[[sv:Ammoniak]]
[[ta:அமோனியா]]
[[th:แอมโมเนีย]]
[[vi:Amoniac]]
[[tr:Amonyak]]
[[uk:Аміак]]
[[zh:氨]]