Haber process
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/* The process */ transferring info from Ammonia and engineering aspects
The '''Haber process''', also called the '''Haber–Bosch process''', is the [[nitrogen fixation]] reaction of [[nitrogen]] and [[hydrogen]], over an iron [[substrate (chemistry)|substrate]], to produce [[ammonia]].<ref name="enriching">''Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production'' by Vaclav Smil (2001) ISBN 0-262-19449-X</ref><ref>''Fertilizer Industry: Processes, Pollution Control and Energy Conservation'' by Marshall Sittig (1979) Noyes Data Corp., N.J. ISBN 0-8155-0734-8</ref><ref>"Heterogeneous Catalysts: A study Guide"</ref> The Haber process is important because ammonia is difficult to produce on an industrial scale and the fertilizer generated from the ammonia is responsible for sustaining one-third of the Earth's population.<ref>{{cite book |last=Wolfe |first=David W. |title=Tales from the underground a natural history of subterranean life |year=2001 |publisher=Perseus Pub |location=Cambridge, Mass |isbn=0738201286 |oclc= 46984480 }}.</ref> Even though 78.1% of the [[air]] we breathe is [[nitrogen]], the gas is relatively unreactive because nitrogen molecules are held together by strong [[Chemical bond|triple bonds]]. It was not until the early 20th century that this method was developed to harness the atmospheric abundance of [[nitrogen]] to create [[ammonia]], which can then be [[redox|oxidized]] to make the [[nitrate]]s and [[nitrite]]s essential for the production of [[nitrate]] [[fertilizers|fertilizer]] and [[munitions]].
==History==
The process was first patented by [[List of famous Germany people|German]] [[List of chemists|chemist]] [[Fritz Haber]]. In [[1910]] [[Carl Bosch]], while working for [[Germany|German]] chemical company [[BASF]], successfully commercialized the process and secured further patents. Haber and Bosch were later awarded [[Nobel Prize|Nobel prizes]], in 1918 and 1931 respectively, for their work in overcoming the chemical and engineering problems posed by the use of large-scale high-pressure technology.
Ammonia was first manufactured using the Haber process on an industrial scale in Germany during [[World War I]], to meet the high demand for [[ammonium nitrate]] (for use in explosives) at a time when supply of [[Sodium nitrate|Chile saltpetre]] from [[Chile]] could not be guaranteed because this industry was then almost 100% in British hands. It has been suggested that without this process, Germany would not have fought in the war<ref>[http://query.nytimes.com/mem/archive-free/pdf?_r=1&res=9807EEDA133BEE32A25750C0A9649C946195D6CF&oref=slogin]</ref>.
==The process==
The bulk of the chemical technology consists of isolating [[hydrogen]] from [[methane]] ([[natural gas]]) using [[heterogeneous catalysis]] and then reacting it with atmospheric nitrogen.
=== Synthesis gas preparation===
First, the methane is cleaned, mainly to remove [[sulfur]] impurities that would poison the catalysts.
The clean methane is then reacted with [[steam]] over a catalyst of [[nickel oxide]]. This is called [[steam reforming]]:
:CH<sub>4</sub> + H<sub>2</sub>O → CO + 3H<sub>2</sub>
Secondary reforming then takes place with the addition of air to convert the methane that did not react during steam reforming.
:2CH<sub>4</sub> + O<sub>2</sub> → 2CO + 4H<sub>2</sub>
:CH<sub>4</sub> + 2O<sub>2</sub> → CO<sub>2</sub> + 2H<sub>2</sub>O
Then two shift reactions convert CO to CO<sub>2</sub> by reaction with steam.
The gas mixture is now passed into a methanator, which converts any remaining CO<sub>2</sub> into methane for recycling:
: CO<sub>2</sub> + 4H<sub>2</sub> → CH<sub>4</sub> + 2H<sub>2</sub>O
===Ammonia synthesis===
The final stage is the synthesis of ammonia using [[magnetite]], iron oxide, as the catalyst:
:N<sub>2</sub>(g) + 3H<sub>2</sub>(g) → 2NH<sub>3</sub>(g), ΔH<sup>o</sup> = -92.4 kJ/[[Mole (unit)|mol]]
This is done at 150 - 250 [[Atmosphere (unit)|atmospheres]] (atm) and between 300 and 550 °C, passing the gases over four beds of catalyst, with cooling between each pass to maintain a reasonable [[equilibrium constant]]. On each pass only about 15% conversion occurs, but any unreacted gases are recycled, so that eventually an overall conversion of 98% can be achieved.
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>
==Reaction rate and equilibrium==
There are two opposing considerations in this synthesis: the position of the equilibrium and the [[reaction rate|rate of reaction]]. At room temperature, the reaction is slow and the obvious solution is to raise the temperature. This may increase the rate of the reaction but, since the reaction is [[exothermic reaction|exothermic]], it also has the effect, according to [[Le Chatelier's Principle]], of favouring the reverse reaction and thus reducing [[equilibrium constant]], given by:
: <math>K_\mathrm{eq} = \mathrm{\frac{[NH_3]^2}{[N_2][H_2]^3}}</math>
{| align="right" border="3" class="wikitable"
|+Variation in K<sub>eq</sub> for the Equilibrium<br>
N<sub>2</sub> (g) + 3H<sub>2</sub> (g) ↔ 2NH<sub>3</sub> (g)<br>
as a Function of Temperature<ref>Chemistry the Central Science" Ninth Ed., by: Brown, Lemay, Bursten, 2003, ISBN 0-13-038168-3</ref>
|-
! Temperature (°C)
! K<sub>eq</sub>
|-
|align="center" | 300
| 4.34 x 10<sup>–3</sup>
|-
|align="center" | 400
| 1.64 x 10<sup>–4</sup>
|-
|align="center" | 450
| 4.51 x 10<sup>–5</sup>
|-
|align="center" | 500
| 1.45 x 10<sup>–5</sup>
|-
|align="center" | 550
| 5.38 x 10<sup>–6</sup>
|-
|align="center" | 600
| 2.25 x 10<sup>–6</sup>
|}
As the temperature increases, the [[equilibrium]] is shifted and hence, the constant drops dramatically according to the [[van't Hoff equation]]. Thus one might suppose that a low temperature is to be used and some other means to increase rate. However, the catalyst itself requires a temperature of at least 400 °C to be efficient.
[[Pressure]] is the obvious choice to favour the forward reaction because there are 4 moles of reactant for every 2 moles of product (see [[entropy]]), and the pressure used (around 200 atm) alters the equilibrium concentrations to give a profitable yield.
Economically, though, pressure is an expensive commodity. Pipes and reaction vessels need to be strengthened, valves more rigorous, and there are safety considerations of working at 200 atm. In addition, running pumps and compressors takes considerable energy. Thus the compromise used gives a single pass yield of around 15%.
Another way to increase the yield of the reaction would be to remove the product (i.e. ammonia gas) from the system. In practice, gaseous ammonia is not removed from the reactor itself, since the temperature is too high; but it is removed from the equilibrium mixture of gases leaving the reaction vessel. The hot gases are cooled enough, whilst maintaining a high pressure, for the ammonia to condense and be removed as liquid. Unreacted hydrogen and nitrogen gases are then returned to the reaction vessel to undergo further reaction.
==Catalysts==
The [[catalyst]] has no effect on the position of [[chemical equilibrium]]; rather, it provides an alternative pathway with lower [[activation energy]] and hence increases the reaction rate, while remaining chemically unchanged at the end of the reaction. The first Haber–Bosch reaction chambers used [[osmium]] and [[uranium]] catalysts. However, today a much less expensive [[iron]] catalyst is used almost exclusively.
In industrial practice, the iron catalyst is prepared by exposing a mass of [[magnetite]], an iron oxide, to the hot hydrogen feedstock. This reduces some of the magnetite to metallic iron, removing [[oxygen]] in the process. However, the catalyst maintains most of its bulk volume during the reduction, and so the result is a highly porous material whose large surface area aids its effectiveness as a catalyst. Other minor components of the catalyst include [[calcium]] and [[aluminium oxide]]s, which support the porous iron catalyst and help it maintain its surface area over time, and [[potassium]], which increases the [[electron]] density of the catalyst and so improves its reactivity.
The [[reaction mechanism]], involving the heterogeneous catalyst, is believed to be as follows:
# N<sub>2</sub>(g) → N<sub>2</sub>(adsorbed)
#''' N<sub>2</sub>(adsorbed) → 2N(adsorbed)'''
# H<sub>2</sub>(g) → H<sub>2</sub>(adsorbed)
# H<sub>2</sub>(adsorbed) → 2H(adsorbed)
# N(adsorbed) + 3H(adsorbed)→ NH<sub>3</sub>(adsorbed)
# NH<sub>3</sub>(adsorbed) → NH<sub>3</sub>(g)
Reaction 5 occurs in three steps, forming NH, NH<sub>2</sub>, and then NH<sub>3</sub>. Experimental evidence points to reaction 2 as being the slow, [[rate-determining step]].
A major contributor to the elucidation of this mechanism is [[Gerhard Ertl]].<ref>''Interaction of nitrogen with iron surfaces : I. Fe(100) and Fe(111)'' Journal of Catalysis, Volume 49, Issue 1, July '''1977''', Pages 18-41 F. Bozso, G. Ertl, M. Grunze and M. Weiss {{DOI|10.1016/0021-9517(77)90237-8}}</ref><ref>The structure of atomic nitrogen adsorbed on Fe(100) Surface Science, Volume 123, Issue 1, 1 December '''1982''', Pages 129-140 R. Imbihl, R. J. Behm, G. Ertl and W. Moritz {{DOI|10.1016/0039-6028(82)90135-2}}</ref><ref>''Kinetics of nitrogen adsorption on Fe(111)'' Surface Science, Volume 114, Issues 2-3, 1 February '''1982''', Pages 515-526 G. Ertl, S. B. Lee and M. Weiss {{DOI|10.1016/0039-6028(82)90702-6}}</ref><ref>''Primary steps in catalytic synthesis of ammonia'' G. Ertl Journal of Vacuum Science & Technology A: Vacuum, Surfaces, and Films -- April 1, '''1983''' -- Volume 1, Issue 2, pp. 1247-1253 {{DOI|10.1116/1.572299}}</ref>
==Economic and environmental aspects==
The Haber process now produces 100 million tons of nitrogen [[fertilizers|fertilizer]] per year, mostly in the form of anhydrous [[ammonia]], [[ammonium nitrate]], and [[urea]]. 3-5% of world natural gas production is consumed in the Haber process (~1-2% of the world's annual energy supply)<ref name="enriching" />,<ref>{{cite web|url=http://www.eia.doe.gov/oiaf/ieo/nat_gas.html | title=International Energy Outlook 2007}}</ref>,<ref>http://www.fertilizer.org/ifa/statistics/indicators/ind_reserves.asp}}</ref>,<ref>{{cite web|url=http://www.sciencemag.org/cgi/content/full/297/5587/1654|title=Science, 6 September 2002: Vol. 297. no. 5587, pp. 1654 - 1655 DOI: 10.1126/science.1076659}}</ref>. That fertilizer is responsible for sustaining one-third of the Earth's population, as well as various deleterious environmental consequences.<ref>{{cite book |last=Wolfe |first=David W. |title=Tales from the underground a natural history of subterranean life |year=2001 |publisher=Perseus Pub |location=Cambridge, Mass |isbn=0738201286 |oclc= 46984480 }}.</ref> Generation of hydrogen using electrolysis of water, using renewable energy, is not currently competitive cost-wise with hydrogen from fossil fuels, such as natural gas, and is responsible for only 4% of current hydrogen production.
Notably, the rise of this industrial process led to the "Nitrate Crisis" in [[Chile]], when the British industrials left the country -- since the natural nitrate mines were no longer profitable -- closing the mines and leaving a large unemployed Chilean population behind.
==See also==
*[[Chemical kinetics]]
*[[Reaction rate]]
*[[Rate equation]]
==References==
{{reflist}}
==External links==
*[http://www.wisegeek.com/what-is-the-haber-bosch-process.htm What is the Haber-Bosch Process?]
*[http://www.idsia.ch/~juergen/haberbosch.html Haber-Bosch process]
*[http://www.efma.org/Publications/10anniversary/Section02.asp Fertilizer,agriculture and the production of food]
*[http://www.britannica.com/nobelprize/article-9038696 Britannica guide to Nobel Prizes: Fritz Haber]
*[http://nobelprize.org/nobel_prizes/chemistry/laureates/1918/haber-bio.html Nobel e-Museum - Biography of Fritz Haber]
*[http://www.ausetute.com.au/haberpro.html Uses and Production of Ammonia]
*”CIEC Catalysis”, [http://www.uyseg.org/catalysis/pages/cat_frames.html]
*[http://www.ias.ac.in/resonance/Sept2002/pdf/Sept2002p69-77.pdf Haber Process for Ammonia Synthesis]
* {{cite web | author = Jim Clark | title = The Haber Process | url = http://www.chemguide.co.uk/physical/equilibria/haber.html | date = 2002 | work = [http://www.chemguide.co.uk Chemguide]}}
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