The Haber process, also called the Haber–Bosch process, is the industrial implementation of the reaction of nitrogen gas and hydrogen gas. It is the main industrial route to ammonia:
N2 + 3 H2 → 2 NH3 (ΔH = −92.4 kJ·mol−1)
Nitrogen is a critical limiting mineral nutrient in plant growth. Carbon and oxygen are also critical, but are easily obtained by plants from soil and air. Even though air is 78% nitrogen, atmospheric nitrogen is nutritionally unavailable because nitrogen molecules are held together by strong triple bonds. Nitrogen must be ‘fixed’, i.e. converted into some bioavailable form, through natural or man-made processes. It was not until the early 20th century that Fritz Haber developed the first practical process to convert atmospheric nitrogen to ammonia, which is nutritionally available. Prior to the discovery of the Haber process, ammonia had been difficult to produce on an industrial scale.
Fertilizer generated from ammonia produced by the Haber process is estimated to be responsible for sustaining one-third of the Earth’s population.[6] 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.
History
Main article: History of the Haber process
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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.[8] 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.[3][9] 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/day the following year.[10] 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 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 important to the German war effort.
The process
This conversion is typically conducted at 15–25 MPa (2,200–3,600 psi) or 150–250 bar and between 300–550 °C (572–1022 °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 absolute pressures of about 2.5–3.5 MPa (360–510 psi) or 25–35 bar, and the ammonia synthesis loop operates at absolute pressures ranging from[clarification needed] 6–18 MPa (870–2,600 psi) or 60–180 bar, depending upon which proprietary process design is used.
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.
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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 atm units) becomes 1 around 150° or 200°C. (See Le Chatelier’s principle.)
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.
Catalysts[edit source | editbeta]
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).
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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 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:
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.
A major contributor to the elucidation of this mechanism is Gerhard Ertl.
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 the world’s natural gas production is consumed in the Haber process (~1–2% of the world’s annual energy supply).That fertilizer is responsible for sustaining one-third of the Earth’s population, but results in various deleterious environmental consequences.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.
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