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 twentieth century. The process converts atmospheric nitrogen (N2) to ammonia (NH3) by a reaction with hydrogen (H2) using a metal catalyst under high temperatures and pressures:
- N2 + 3 H2 → 2 NH3 (ΔH = −92.4 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.
Although the Haber process is mainly used to produce fertilizer today, during 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.
Throughout the nineteenth century the demand for nitrates and ammonia for use as fertilizers and industrial feedstocks had been steadily increasing; however, the main source remained the mining of niter deposits. By the start of the twentieth century it was being predicted that these reserves would be unable to satisfy future demand and research into new potential sources of ammonia became ever-more important. The most obvious source was atmospheric nitrogen (N2), which makes up nearly 80% of the air, however N2 is exceptionally stable and will not readily react with other chemicals. Converting N2 into ammonia therefore posed a chemical challenge which occupied the efforts of chemists across the world.
Haber together with his assistant Robert Le Rossignol developed the high-pressure devices and catalysts used 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 assigned Carl Bosch the task of scaling up Haber's tabletop machine to industrial-level production. He succeeded in this process 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 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, production reaching 20 tonnes per day the following year. During World War I, the synthetic ammonia was used for the production of nitric acid, a precursor to munitions. The Allies had access to large amounts of sodium nitrate deposits in Chile (so called "Chile saltpetre") that belonged almost totally to British industries. As Germany lacked access to such readily available natural resources, the Haber process proved essential to the continued German war effort.
This conversion is typically conducted at 15–25 MPa (2,200–3,600 psi) or 150–250 bar and between 400–500 °C (752–932 °F), 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.
The steam reforming, shift conversion, carbon dioxide removal, and methanation steps each operate at pressures of about 2.5–3.5 MPa (360–510 psi) or 25–35 bar, and the ammonia synthesis loop operates at pressures ranging from 6–18 MPa (870–2,610 psi) or 60–180 bar, 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 air, which is deoxygenated by the combustion of 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 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 atm units) becomes 1 around 150° or 200 °C. (See Le Chatelier's principle.)
|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|
Above this temperature, the equilibrium quickly becomes quite unfavourable 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 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.
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 researcher 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 on graphite.
In industrial practice, the iron catalyst is obtained from finely ground iron powder, which in turn 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), 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 immune to reduction by the hydrogen.
- 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 broken.
Economic and environmental aspects
When it was first invented, the Haber process needed to compete against another industrial processes, the Cyanamide process. However, the Cyanamide process consumed large amounts of electrical power and was more labor-intensive than the Haber process.:137–143
The Haber process now produces 450 million tonnes (440,000,000 long tons; 500,000,000 short tons) of nitrogen fertilizer per year, mostly in the form of anhydrous ammonia, ammonium nitrate, and urea. 3–5% of the world's natural gas production is consumed in the Haber process (~1–2% of the world's annual energy supply). In combination with pesticides, these fertilizers have quadrupled the productivity of agricultural land:
- 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 required today.
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 80% of the nitrogen found in human tissues originated from the Haber-Bosch process. Since nitrogen use efficiency is typically less than 50%, our heavy use of industrial nitrogen fixation is severely disruptive to our biological habitat.
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