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-==Synthesis and production==
-{{about|section=true|industrial synthesis|synthesis in certain organisms|#Biosynthesis}}
-{{See also|Ammonia production}}
-[[File:Production of ammonia.svg|thumb|upright=1.35|right|Production trend of ammonia between 1947 and 2007]]
-Ammonia is one of the most produced inorganic chemicals, with global production reported at 176 million tonnes in 2014.<ref name="USGS2016"/> China accounted for 32.6% of that, followed by Russia at 8.1%, India at 7.6%, and the United States at 6.4%.<ref name="USGS2016"/>
-Before the start of [[World War I]], most ammonia was obtained by the [[dry distillation]]<ref>{{cite news|accessdate=7 July 2009|url=http://nobelprize.org/chemistry/laureates/1918/press.html |title=Nobel Prize in Chemistry (1918) – Haber-Bosch process}}</ref> of nitrogenous vegetable and animal waste products, including [[camel]] [[manure|dung]], where it was [[distillation|distilled]] by the reduction of [[nitrous acid]] and [[nitrite]]s with hydrogen; in addition, it was produced by the distillation of [[coal]], and also by the decomposition of ammonium salts by [[alkaline]] hydroxides<ref>{{cite news|accessdate=7 July 2009|title=Chemistry of the Group 2 Elements – Be, Mg, Ca, Sr, Ba, Ra|url=http://www.bbc.co.uk/dna/h2g2/A1002934|publisher= BBC.co.uk}}</ref> such as [[calcium oxide|quicklime]], the salt most generally used being the chloride ([[sal ammoniac]]) thus:{{sfn|Chisholm|1911|p=861}}
-:2 NH<sub>4</sub>Cl + 2 CaO → [[Calcium chloride|CaCl<sub>2</sub>]] + Ca(OH)<sub>2</sub> + 2 NH<sub>3([[Gas|g]])</sub>
-For small scale laboratory synthesis, one can heat [[urea]] and [[calcium hydroxide]]:
-:(NH2)<sub>2</sub>CO + Ca(OH)<sub>2</sub> → CaCO<sub>3</sub> + 2 NH<sub>3</sub>
-===Haber-Bosch Process===
-{{See also|Haber-Bosch Process}}
-Modern ammonia production mostly uses the [[Haber process|Haber–Bosch process]], reacting hydrogen (H<sub>2</sub>) and nitrogen (N<sub>2</sub>) at a moderately-elevated temperature (450&nbsp;°C) and high pressure ({{convert|100|atm}}):<ref>Atkins, P.W.; Overton, T.L.; Rourke, J.P.; Weller, M.T. and Armstrong, F.A. (2010) [https://docs.google.com/file/d/0B4Ka5HSSrR_yeENFS0hvbjdUR00/view ''Shriver and Atkins Inorganic Chemistry'']. 5th Edi. W. H. Freeman and Company, New York. p. 383. {{ISBN|978-1-42-921820-7}}</ref>
-:3 H<sub>2</sub> + N<sub>2</sub> → 2 NH<sub>3([[Gas|g]])</sub>
-This reaction is both exothermic and results in decreased entropy, meaning that the [[Gibbs free energy|reaction is favoured]] at lower temperatures<ref>See {{harv|Clark|2013}}: "The forward reaction (the production of ammonia) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. In order to get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible".</ref> and higher pressures.<ref>See {{harv|Clark|2013}}: "Notice that there are 4 molecules on the left-hand side of the equation, but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high".</ref> This makes it difficult and expensive to achieve, as lower temperatures result in slower [[reaction kinetics]] (hence a slower [[reaction rate]])<ref>See {{harv|Clark|2013}}: "However, 400 - 450°C isn't a low temperature! Rate considerations: The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium".</ref> and high pressure requires high-strength pressure vessels<ref>See {{harv|Clark|2013}}: "Rate considerations: Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction. Economic considerations: Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built".</ref> that aren't weakened by [[hydrogen embrittlement]]. In addition, [[diatomic]] nitrogen is bound together by an exceptionally strong [[triple bond]], which makes it rather inert. Both the yield and efficiency of the Haber-Bosch Process are low, meaning that ammonia produced must be continuously separated and extracted for the reaction to proceed at an appreciable pace. Combined with the energy needed to [[Hydrogen production|produce hydrogen]]{{refn|group=note|Hydrogen required for ammonia synthesis is most often produced through [[gasification]] of cabon-containing material, mostly natural gas, but other potential carbon sources include coal, petroleum, peat, biomass, or waste. As of 2012, the global production of ammonia produced from natural gas using the steam reforming process was 72 percent.<ref>{{Cite news|url=http://ietd.iipnetwork.org/content/ammonia|title=Ammonia|date=2013-04-30|work=Industrial Efficiency Technology & Measures|access-date=2018-04-06|language=en}}</ref> Hydrogen can also be produced from water and electricity using [[Electrolysis of water|electrolysis]]: at one time, most of Europe's ammonia was produced from the Hydro plant at [[Vemork]]. Other possibilities include [[biological hydrogen production]] or [[photolysis]], but at present, [[steam reforming]] of natural gas is the most economical means of mass-producing hydrogen.}} and purified atmospheric nitrogen, ammonia production is a very energy-intensive process, consuming 1 to 2% of global energy, 3% of global carbon emissions<ref>{{cite web |url=https://phys.org/news/2018-07-electrochemically-produced-ammonia-revolutionize-food-production.html |title= Electrochemically-produced ammonia could revolutionize food production |author= [[Lehigh University]] |date= 2018-07-09 |language= en |access-date= 2018-12-15 |quote="Ammonia manufacturing consumes 1 to 2% of total global energy and is responsible for approximately 3% of global carbon dioxide emissions."}}</ref>, and 3 to 5% of natural gas consumption<ref>{{cite web |url=https://www.ornl.gov/content/physical-catalyst-electrolysis-nitrogen-ammonia |title=A physical catalyst for the electrolysis of nitrogen to ammonia |last=Song |first=Yang |last2=Hensley |first2=Dale |last3=Bonnesen |first3=Peter |last4=Liang |first4=Liango |last5=Huang |first5=Jingsong |last6=Baddorf |first6=Arthur |last7=Tschaplinski |first7=Timothy  |last8=Engle |first8=Nancy |last9=Wu |first9=Zili |last10=Cullen |first10=David |last11=Meyer |first11=Harry III |last12=Sumpter |first12=Bobby |last13=Rondinone |first13=Adam |date=2018-05-02 |year=2018 |publisher=Oak Ridge National Laboratory |language= en |doi=10.1126/sciadv.1700336 |access-date= 2018-12-15 |quote= "Ammonia synthesis consumes 3 to 5% of the world’s natural gas, making it a significant contributor to greenhouse gas emissions."}}</ref>.
-*{{cite web |url=https://chemguide.co.uk/physical/equilibria/haber.html |title=THE HABER PROCESS |last=Clark |first=Jim |date= April 2013 |orig-year= 2002 |language= en |access-date= 15 Dec 2018 |ref= {{harvid|Clark|2013}} }}
-==Notes==
-{{Reflist|group=note}}
-==References==
-{{Reflist}}