Haber process

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Fritz Haber, 1918

The Haber process, also called the Haber–Bosch process, is the nitrogen fixation reaction of nitrogen gas and hydrogen gas, over an enriched iron or ruthenium catalyst, which is used to industrially produce ammonia.[1][2][3][4]

Despite the fact that 78.1% of the air we breathe is nitrogen, the gas is relatively unavailable because it is so unreactive: nitrogen molecules are held together by strong triple bonds. It was not until the early 20th century that Fritz Haber developed the first practical process to convert atmospheric nitrogen to ammonia. Prior to the discovery of the Haber process, ammonia had been difficult to produce on an industrial scale.

The Haber process is important today because the fertilizer generated from ammonia is responsible for sustaining one-third of the Earth's population.[5] It is estimated that half of the protein within human beings is made of nitrogen that was originally fixed by this process, the remainder was produced by nitrogen fixing bacteria and archaea.[6]

Contents

History

Early in the twentieth century, several chemists tried to make ammonia from atmospheric nitrogen. German chemist Fritz Haber discovered a process that is still used today. Robert Le Rossignol was instrumental in the development of the high-pressure devices used in the Haber process.[7] They demonstrated their process in the summer of 1909 by producing ammonia from air drop by drop, at the rate of about 125 ml (4 US fl oz) per hour. The process was purchased by the German chemical company BASF, which assigned Carl Bosch the task of scaling up Haber's tabletop machine to industrial-level production.[2][8] Haber and Bosch were later awarded Nobel prizes, in 1918 and 1931 respectively, for their work in overcoming the chemical and engineering problems posed by the use of large-scale, continuous-flow, high-pressure technology.

Ammonia was first manufactured using the Haber process on an industrial scale in 1913 in BASF's Oppau plant in Germany. During World War I, the synthetic ammonia was utilized for the production of nitric acid, a precursor to munitions. The Allies had access to large amounts of saltpeter deposits in Chile that belonged almost totally to British industries. The Habe process proved important to the German war effort.[9]

The process

A historical (1921) high-pressure steel reactor for production of ammonia via Bosch-Haber process at the BASF plant in Ludwigshafen, Germany.

The Haber process is the main industrial route for the synthesis of ammonia:[10]

N2 + 3 H2 ⇌ 2 NH3   (ΔH = −92.22 kJ·mol−1)

This conversion is typically conducted at 15–25 MPa (150–250 bar) and between 300 and 550 °C, as the gases are passed over four beds of catalyst, with cooling between each pass so as to maintain a reasonable equilibrium constant. On each pass only about 15% conversion occurs, but any unreacted gases are recycled, and eventually an overall conversion of 97% is achieved.[10]

The steam reforming, shift conversion, carbon dioxide removal, and methanation steps each operate at absolute pressures of about 2.5–3.5 MPa (25–35 bar), and the ammonia synthesis loop operates at absolute pressures ranging from 6–18 MPa (59–178 atm), depending upon which proprietary design is used.[10]

Sources of hydrogen

The major source is methane from natural gas. The conversion, steam reforming, is conducted with air, which is deoxygenated by the combusting natural gas. Originally Bosch obtained hydrogen by the electrolysis of water.

Reaction rate and equilibrium

Two opposing considerations are relevant to this synthesis: the position of the equilibrium and the 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, it also has the effect, according to Le Chatelier's principle, of favouring the reverse reaction and thus reducing the amount of product.

Variation in Keq for the equilibrium
N2 (g) + 3H2 (g) is in equilibrium with 2NH3 (g)
as a function of temperature[11]
Temperature (°C) Keq
300 4.34 x 10−3
400 1.64 x 10−4
450 4.51 x 10−5
500 1.45 x 10−5
550 5.38 x 10−6
600 2.25 x 10−6

As the temperature increases, the equilibrium is shifted and hence, the amount of product 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 most popular catalysts are based on iron promoted with K2O, CaO, SiO2, and Al2O3. The original Haber–Bosch reaction chambers used osmium as catalysts. However, under Bosch's direction in 1909, the BASF researcher Alwin Mittasch discovered a much less expensive iron-based catalyst, which is still used today. Part of the industrial production utilizes ruthenium rather than an iron-based catalysts (the KAAP process). Ruthenium forms more active catalysts that allows milder operating pressures. Such catalysts are prepared by decomposition of triruthenium dodecacarbonyl on graphite.[10]

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 oxides, 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 activity.

The reaction mechanism, involving the heterogeneous catalyst, is believed to involve the following steps:

  1. N2 (g) → N2 (adsorbed)
  2. N2 (adsorbed) → 2 N (adsorbed)
  3. H2(g) → H2 (adsorbed)
  4. H2 (adsorbed) → 2 H (adsorbed)
  5. N (adsorbed) + 3 H(adsorbed)→ NH3 (adsorbed)
  6. NH3 (adsorbed) → NH3 (g)

Reaction 5 occurs in three steps, forming NH, NH2, and then NH3. 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.[12]

Economic and environmental aspects

The Haber process now produces 500 million tons (453 billion kilograms) of nitrogen 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).[1][13][14][15] That fertilizer is responsible for sustaining one-third of the Earth's population, but results in various deleterious environmental consequences.[2][5] Hydrogen production using electrolysis of water powered by renewable energy is not yet competitive with hydrogen from fossil fuels, such as natural gas. As of 2007, only 5% of hydrogen is produced by electrolysis.[16] Notably, the rise of the Haber industrial process led to the "Nitrate Crisis" in Chile when the natural nitrate mines were no longer profitable and were closed, leaving a large unemployed Chilean population behind.

See also

References

  1. 1.0 1.1 Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production by Vaclav Smil (2001) ISBN 0-262-19449-X
  2. 2.0 2.1 2.2 Hager, Thomas (2008). The Alchemy of Air. Harmony Books, New York. ISBN 978-0-307-35178-4.
  3. Fertilizer Industry: Processes, Pollution Control and Energy Conservation by Marshall Sittig (1979) Noyes Data Corp., N.J. ISBN 0-8155-0734-8
  4. "Heterogeneous Catalysts: A study Guide"
  5. 5.0 5.1 Script error
  6. BBC: Discovery - Can Chemistry Save The World? - 2. Fixing the Nitrogen Fix
  7. "Robert Le Rossignol, 1884–1976: Professional Chemist", ChemUCL Newsletter (UCL Department of Chemistry): 8, 2009, http://www.ucl.ac.uk/chemistry/alumni/documents/A5booklet_020909.pdf
  8. US Pat 990191
  9. "Nobel Award to Haber". New York Times. 3 February 1920. http://query.nytimes.com/mem/archive-free/pdf?_r=1&res=9807EEDA133BEE32A25750C0A9649C946195D6CF&oref=slogin. Retrieved 11 October 2010.
  10. 10.0 10.1 10.2 10.3 Max Appl "Ammonia" in Ullmann's Encyclopedia of Industrial Chemistry 2006 Wiley-VCH, Weinheim. doi:10.1002/14356007.a02_143.pub2
  11. Chemistry the Central Science" Ninth Ed., by: Brown, Lemay, Bursten, 2003, ISBN 0-13-038168-3
  12. Script error. Script error. Script error. Script error
  13. "International Energy Outlook 2007". http://www.eia.doe.gov/oiaf/ieo/nat_gas.html.
  14. "?". http://www.fertilizer.org/ifa/statistics/indicators/ind_reserves.asp.
  15. Script error
  16. Peter Häussinger, Reiner Lohmüller, Allan M. Watson “Hydrogen” Ullmann's Encyclopedia of Industrial Chemistry, 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a13_297

External links

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