My watch list  

Haber process

The Haber process (also known as Haber–Bosch process) is the reaction of nitrogen and hydrogen, over an iron-substrate, to produce ammonia.[1][2][3] The Haber process is important because ammonia is difficult to produce, on an industrial scale. Even though 78.1% of the air we breathe is nitrogen, the gas is relatively inert due to the strength of the triple bond that keeps the molecule together. It was not until the start of the twentieth century that this method was developed to harness the atmospheric abundance of nitrogen to create ammonia, which can then be oxidised to make the nitrates and nitrites essential for the production of nitrate fertilizer and munitions.



In the Haber Process, nitrogen (N2) and hydrogen (H2) gases are reacted over an iron catalyst (Fe3+) in which aluminium oxide (Al2O3) and potassium oxide (K2O) are used as promoters. The reaction is carried out under conditions of 250 atmospheres (atm), 450-500 °C; resulting in a yield of 10-20%:

N2(g) + 3H2(g) → 2NH3(g) ΔHo = -92.4 kJ/mol

(Where ΔHo is the standard heat of reaction or standard enthalpy change)

These conditions are chosen due to the high reaction rate which they foster despite the poor relative amount of ammonia produced.


The process was first patented by Fritz Haber. In 1910 Carl Bosch, while working for chemical company BASF, successfully commercialized the process and secured further patents. 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 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 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 almost certainly have run out of explosives by 1916, thereby ending the war.

The process

The bulk of the chemical technology consists in getting the hydrogen from methane or natural gas using heterogeneous catalysis and then reacting it with the atmospheric nitrogen.

Synthesis gas preparation

First, the methane is cleaned, mainly to remove sulphur impurities that would poison the catalysts. This is done by turning sulphur into hydrogen sulphide:

CH3SH + H2 → CH4 + H2S

and then reacting with zinc oxide to form zinc sulphide:

H2S + ZnO → ZnS + H2O

The clean methane is then reacted with steam over a catalyst of nickel oxide. This is called steam reforming and occurs in two steps:

First step: (one mole of methane in)

CH4 + H2O → CO + 3H2 (3 moles of hydrogen out)
CO + H2O → CO2 + H2 (1 extra mole of hydrogen out)

Note that 4 moles of hydrogen are produced per mole of methane

Secondary reforming then takes place with the addition of air:

2 H2 + O2 + N2 → 2H2O + N2

This now gives a ratio of nitrogen to hydrogen of 1:5

Then occur two’’shifts’’ which take CO to CO2 again by reaction with steam, one at high temperature, then one at low temperature:

CO + H2O → CO2 + H2 high temperature 1:6
→ the catalyst here is a mixture of iron, chromium and copper
CO + H2O → CO2 + H2 low temperature 1:7
→ the catalyst here is a mixture of copper, zinc and aluminum

The removal of carbon dioxide is easily done by reaction with potassium carbonate.

K2CO3 + H2O + CO2→ 2KHCO3

The gas mixture is now passed into a methanator which converts any remaining CO2 into methane for recycling:

CO2 + 4H2 → CH4 + 2H2O 1:3

We now have a gas mixture containing nitrogen and hydrogen in the correct ratio of 1:3. This is called synthesis gas.

Ammonia synthesis

The final stage is the crucial synthesis of ammonia using promoted magnetite, iron oxide, as the catalyst:

N2(g) + 3H2(g) → 2NH3(g), ΔHo = -92.4 kJ/mol

This is done at 100 - 250 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 will be recycled, so that eventually an overall yield of 98% can be achieved.

Reaction Rate and Equilibrium

There are two opposing considerations in 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 equilibrium constant, given by:

K_\mathrm{eq} = \mathrm{\frac{[NH_3]^2}{[N_2][H_2]^3}}

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, 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.


The catalyst has no effect on the position of 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 aluminum 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 reactivity.

The reaction mechanism, involving the heterogeneous catalyst, is believed to be as follows:

  1. N2(g) → N2(adsorbed)
  2. N2(adsorbed) → 2N(adsorbed)
  3. H2(g) → H2(adsorbed)
  4. H2(adsorbed) → 2H(adsorbed)
  5. N(adsorbed) + 3H(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.[4][5][6][7]

Economic and environmental aspects

The Haber process now produces 100 million tons of nitrogen fertilizer per year, mostly in the form of anhydrous ammonia, ammonium nitrate, and urea. 0.75% of the world's annual energy supply is consumed in the Haber process (3.35% of world natural gas production is used for ammonia production,[8][9][1] and natural gas represents 22% of world energy production.[10] See also[11] for rough estimate of 1% of energy production.) That fertilizer is responsible for sustaining one-third of the Earth's population, as well as various deleterious environmental consequences.[12] 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


  1. ^ a b Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production by Vaclav Smil (2001) ISBN 0-262-19449-X
  2. ^ Fertilizer Industry: Processes, Pollution Control and Energy Conservation by Marshall Sittig (1979) Noyes Data Corp., N.J. ISBN 0-8155-0734-8
  3. ^ “Heterogeneous Catalysts: A study Guide”
  4. ^ 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
  5. ^ 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
  6. ^ 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
  7. ^ 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
  8. ^ International Energy Outlook 2007.
  9. ^ Why Are Nitrogen Prices So High?.
  10. ^ Highlights.
  11. ^ Science, 6 September 2002: Vol. 297. no. 5587, pp. 1654 - 1655 DOI: 10.1126/science.1076659.
  12. ^ Wolfe, David W. (2001). Tales from the underground a natural history of subterranean life. Cambridge, Mass: Perseus Pub. ISBN 0738201286. OCLC 46984480. .
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Haber_process". A list of authors is available in Wikipedia.
Your browser is not current. Microsoft Internet Explorer 6.0 does not support some functions on Chemie.DE