The Haber process, also called the Haber–Bosch process, is an
artificial nitrogen fixation process and is the main industrial
procedure for the production of ammonia today. It is named after
its inventors, the German chemists
Fritz Haber and Carl Bosch, who
developed it in the first half of the 20th century. The process
converts atmospheric nitrogen (N2) to ammonia (NH3) by a reaction with
hydrogen (H2) using a metal catalyst under high temperatures and
displaystyle ce N2 + 3H2 -> 2NH3
(ΔH° = −91.8 kJ) =>
(ΔH° = −45.8 kJ·mol−1)
Before the development of the Haber process, ammonia had been
difficult to produce on an industrial scale with early
methods such as the
Birkeland–Eyde process and Frank–Caro process
all being highly inefficient.
Haber process is mainly used to produce fertilizer today,
World War I
World War I it provided
Germany with a source of ammonia for
the production of explosives, compensating for the Allied trade
blockade on Chilean saltpeter.
2.1 Sources of hydrogen
Reaction rate and equilibrium
3 Economic and environmental aspects
4 See also
6 External links
Main article: History of the Haber process
Throughout the 19th century the demand for nitrates and ammonia for
use as fertilizers and industrial feedstocks had been steadily
increasing. The main source was mining niter deposits. At the
beginning of the 20th century it was being predicted that these
reserves could not satisfy future demands  and research into new
potential sources of ammonia became more important. The obvious source
was atmospheric nitrogen (N2), comprising nearly 80% of the air,
however N2 is exceptionally stable and will not readily react with
other chemicals. Converting N2 into ammonia posed a challenge for
Haber, with his assistant Robert Le Rossignol, developed the
high-pressure devices and catalysts needed to demonstrate the Haber
process at laboratory scale. 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
Carl Bosch the task of scaling up Haber's tabletop machine to
industrial-level production. He succeeded in 1910. Haber and
Bosch were later awarded Nobel prizes, in 1918 and 1931 respectively,
for their work in overcoming the chemical and engineering problems 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, reaching 20
tonnes per day the following year. During World War I, the
production of munitions required large amounts of nitrate. The Allies
had access to large sodium nitrate deposits in
controlled by British companies.
Germany had no such resources, so the
Haber process proved essential to the German war effort.
Synthetic ammonia from the
Haber process was used for the production
of nitric acid, a precursor to the nitrates used in explosives.
A historical (1921) high-pressure steel reactor for production of
ammonia via the
Haber process is displayed at the Karlsruhe Institute
of Technology, Germany.
scheme of ammoniac reactor
This conversion is typically conducted at 15–25 MPa
(150–250 atm; 2,200–3,600 psi) and between
400–500 °C (752–932 °F), as the gases (nitrogen and
hydrogen) 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.
The steam reforming, shift conversion, carbon dioxide removal, and
methanation steps each operate at pressures of about
2.5–3.5 MPa (25–35 bar; 360–510 psi), and the
ammonia synthesis loop operates at pressures ranging from
6–18 MPa (60–180 bar; 870–2,610 psi), depending
upon which proprietary process is used.
Sources of hydrogen
The major source of hydrogen is methane from natural gas. The
conversion, steam reforming, is conducted with steam in a high
temperature and pressure tube inside a reformer with a nickel
catalyst, separating the carbon and hydrogen atoms in the natural gas.
Reaction rate and equilibrium
Nitrogen (N2) is very unreactive because the molecules are held
together by strong triple bonds. The
Haber process relies on catalysts
that accelerate the scission of this triple bond.
Two opposing considerations are relevant to this synthesis: the
position of the equilibrium and the rate of reaction. At room
temperature, the equilibrium is strongly in favor of ammonia, but the
reaction doesn't proceed at a detectable rate. The obvious solution is
to raise the temperature, but because the reaction is exothermic, the
equilibrium constant (using bar or atm units) becomes 1 around
150–200 °C (302–392 °F). (See Le Châtelier's
Kp(T) for N
2 + 3 H
2 ⇌ 2 NH
4.34 × 10−3
1.64 × 10−4
4.51 × 10−5
1.45 × 10−5
5.38 × 10−6
2.25 × 10−6
Above this temperature, the equilibrium quickly becomes quite
unfavorable at atmospheric pressure, 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 favor the forward reaction because
there are 4 moles of reactant for every 2 moles of product
(see entropy), and the pressure used (15–25 MPa
(150–250 bar; 2,200–3,600 psi)) alters the equilibrium
concentrations to give a profitable yield.
Economically, pressure is an expensive commodity. Pipes, valves, and
reaction vessels need to be strengthened, and there are safety
considerations of working at 20 MPa. 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; 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.[citation
The most popular catalysts are based on iron promoted with K2O, CaO,
SiO2, and Al2O3. The original Haber–Bosch reaction chambers used
osmium as the catalyst, but it was available in extremely small
quantities. Haber noted uranium was almost as effective and easier to
obtain than osmium. Under Bosch's direction in 1909, the BASF
Alwin Mittasch discovered a much less expensive iron-based
catalyst, which is still used today. Some ammonia production utilizes
ruthenium-based catalysts (the KAAP process).
Ruthenium forms more
active catalysts that allows milder operating pressures. Such
catalysts are prepared by decomposition of triruthenium dodecacarbonyl
In industrial practice, the iron catalyst is obtained from finely
ground iron powder, which is usually obtained by reduction of high
purity magnetite (Fe3O4). The pulverized iron metal is burnt
(oxidized) to give magnetite of a defined particle size. The magnetite
particles are then partially reduced, removing some of the oxygen in
the process. The resulting catalyst particles consist of a core of
magnetite, encased in a shell of wüstite (FeO, ferrous oxide), which
in turn is surrounded by an outer shell of iron metal. The catalyst
maintains most of its bulk volume during the reduction, resulting in a
highly porous high surface area material, which enhances its
effectiveness as a catalyst. Other minor components of the catalyst
include calcium and aluminium oxides, which support the iron catalyst
and help it maintain its surface area. These oxides of Ca, Al, K, and
Si are unreactive to reduction by the hydrogen.
The reaction mechanism, involving the heterogeneous catalyst, is
believed to involve the following steps:
N2 (g) → N2 (adsorbed)
N2 (adsorbed) → 2 N (adsorbed)
H2 (g) → H2 (adsorbed)
H2 (adsorbed) → 2 H (adsorbed)
N (adsorbed) + 3 H(adsorbed)→ NH3 (adsorbed)
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. This is not unexpected since the bond broken,
the nitrogen triple bond, is the strongest of the bonds that must be
A major contributor to the elucidation of this mechanism is Gerhard
Economic and environmental aspects
Ammonia production § Sustainable ammonia
Industrial fertilizer plant
When it was first invented, the
Haber process needed to compete
against another industrial process, the Cyanamide process. However,
the Cyanamide process consumed large amounts of electrical power and
was more labor-intensive than the Haber process.:137–143
Haber process now produces 450 million tonnes of nitrogen
fertilizer per year, mostly in the form of anhydrous ammonia, ammonium
nitrate, and urea. Three to five percent of the world's natural gas
production is consumed in the
Haber process (around 1–2% of the
world's annual energy supply). In combination with
pesticides, these fertilizers have quadrupled the productivity of
With average crop yields remaining at the 1900 level the crop harvest
in the year 2000 would have required nearly four times more land and
the cultivated area would have claimed nearly half of all ice-free
continents, rather than under 15% of the total land area that is
Due to its dramatic impact on the human ability to grow food, the
Haber process served as the "detonator of the population explosion",
enabling the global population to increase from 1.6 billion in 1900 to
today's 7 billion. Nearly 50% of the nitrogen found in human
tissues originated from the Haber-Bosch process. Since nitrogen
use efficiency is typically less than 50%, farm runoff from heavy
use of fixed industrial nitrogen disrupts biological habitats.
^ a b c d e Appl, Max (2005), "Ammonia", Ullmann's Encyclopedia of
Industrial Chemistry, Weinheim: Wiley-VCH
^ a b c Smil, Vaclav (2004). Enriching the Earth: Fritz Haber, Carl
Bosch, and the Transformation of World Food Production (1st ed.).
Cambridge, MA: MIT. ISBN 9780262693134.
^ a b c d e Hager, Thomas (2008). The Alchemy of Air: A Jewish genius,
a doomed tycoon, and the scientific discovery that fed the world but
fueled the rise of Hitler (1st ed.). New York, NY: Harmony Books.
^ Sittig, Marshall (1979).
Fertilizer Industry: Processes, Pollution
Control, and Energy Conservation. Park Ridge, NJ: Noyes Data Corp.
^ James, Laylin K. (1993). Nobel Laureates in Chemistry 1901–1992
(3rd ed.). Washington,DC: American Chemical Society. p. 118.
^ Haber, Fritz (2012). Thermodynamik technischer Gasreaktionen (in
German) (1st ed.). Paderborn: Salzwasser Verlag.
^ "Robert Le Rossignol, 1884–1976: Professional Chemist" (PDF),
ChemUCL Newsletter, UCL Department of Chemistry: 8, 2009, archived
from the original (PDF) on 2011-01-13
^ Patent US 990191
^ Philip & Phyllis Morris, "From Fertile Minds" (review) American
^ "Nobel Award to Haber". New York Times. 3 February 1920. Retrieved
11 October 2010.
^ Brown, Theodore L.; LeMay, H. Eugene, Jr; Bursten, Bruce E (2006).
"Table 15.2". Chemistry: The Central Science (10th ed.). Upper Saddle
River, NJ: Pearson. ISBN 0-13-109686-9.
^ Wennerström, Håkan; Lidin, Sven. "Scientific Background on the
Nobel Prize in Chemistry 2007 Chemical Processes on Solid Surfaces"
(PDF). NobelPrize.org. Swedish Academy of Sciences. Retrieved
^ Bozso, F.; Ertl, G.; Grunze, M.; Weiss, M. (1977). "Interaction of
nitrogen with iron surfaces: I. Fe(100) and Fe(111)". J. Catal. 49
(1): 18–41. doi:10.1016/0021-9517(77)90237-8. . Imbihl, R.;
Behm, R. J.; Ertl, G.; Moritz, W. (1982). "The structure of atomic
nitrogen adsorbed on Fe(100)". Surf. Sci. 123 (1): 129–140.
Bibcode:1982SurSc.123..129I. doi:10.1016/0039-6028(82)90135-2. .
Ertl, G.; Lee, S. B.; Weiss, M. (1982). "Kinetics of nitrogen
adsorption on Fe(111)". Surf. Sci. 114 (2–3): 515–526.
Bibcode:1982SurSc.114..515E. doi:10.1016/0039-6028(82)90702-6. .
Ertl, G. (1983). "Primary steps in catalytic synthesis of ammonia". J.
Vac. Sci. Tech. A. 1 (2): 1247–1253. doi:10.1116/1.572299.
^ "International Energy Outlook 2007".
Fertilizer statistics. "?". Archived from the original on
^ Smith, Barry E. (September 2002). "Structure. Nitrogenase reveals
its inner secrets". Science. 297 (5587): 1654–5.
doi:10.1126/science.1076659. PMID 12215632.
^ Smil, Vaclav (2011). "
Nitrogen cycle and world food production"
(PDF). World Agriculture. 2: 9–13.
^ Smil, Vaclav (1999). "Detonator of the population explosion" (PDF).
Nature. 400: 415. Bibcode:1999Natur.400..415S.
^ Erisman, Jan Willem; Sutton, Mark A.; Galloway, James; Klimont,
Zbigniew; Winiwarter, Wilfried (2008-09-28). "How a century of ammonia
synthesis changed the world". Nature Geoscience. 1 (10): 636–639.
^ Oenema, O.; Witzke, H.P.; Klimont, Z.; Lesschen, J.P.; Velthof, G.L.
(2009). "Integrated assessment of promising measures to decrease
nitrogen losses in agriculture in EU-27". Agriculture, Ecosystems and
Environment. 133: 280–288. doi:10.1016/j.agee.2009.04.025.
^ Howarth, R. W. (2008). "Coastal nitrogen pollution: a review of
sources and trends globally and regionally". Harmful Algae. 8:
"The Haber Process". Chemguide.co.uk.
Haber-Bosch process, most important invention of the 20th century,
according to V. Smil, Nature, July 29, 1999, p 415 (by Jürgen
Britannica guide to Nobel Prizes: Fritz Haber
Nobel e-Museum - Biography of Fritz Haber
Fertilizer out of thin air
Uses and Production of Ammonia
Haber Process for
Review of "Between Genius and Genocide: The Tragedy of Fritz Haber,
Father of Chemical Warfare" by Daniel Charles