The Haber Process
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 pressures:
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 WWI it provided Germany with a source of ammonia for the production of explosives, compensating for the Allied trade blockade on Chilean saltpeter.
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. .
Economic Environment Aspects
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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 The 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 agricultural land .
Reaction Rate & 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).