What is Haber bosch process? which catalyst he used?
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The Haber process [1] , also called the
Haber–Bosch process , is an artificial
nitrogen fixation process and is the main industrial procedure for the
production of ammonia today. [2][3] It is named after its inventors, the German chemists Fritz Haber and Carl Bosch , who developed it in the first decade of the 20th century. The process converts atmospheric
nitrogen (N2 ) to ammonia (NH 3) by a reaction with hydrogen (H 2) 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, [4][5][6] 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 Powers ' trade blockade on
Chilean saltpeter.
History
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 [7] 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 chemists globally.
Haber, with his assistant Robert Le Rossignol , developed the high-pressure devices and catalysts needed to demonstrate the Haber process at laboratory scale. [8][9] 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. [5][10] 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. [5]
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.[11] During World War I , the production of
munitions required large amounts of nitrate. The Allies had access to large
sodium nitrate deposits in Chile (Chile saltpetre) controlled by British companies. Germany had no such resources, so the Haber process proved essential to the German war effort. [5][12] Synthetic ammonia from the Haber process was used for the production of nitric acid, a precursor to the nitrates used in explosives.
Process
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. [3]
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. [3]
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 (N 2) 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 principle .)
Kp (T ) for N 2 + 3 H 2 ⇌ 2 NH 3
Haber–Bosch process , is an artificial
nitrogen fixation process and is the main industrial procedure for the
production of ammonia today. [2][3] It is named after its inventors, the German chemists Fritz Haber and Carl Bosch , who developed it in the first decade of the 20th century. The process converts atmospheric
nitrogen (N2 ) to ammonia (NH 3) by a reaction with hydrogen (H 2) 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, [4][5][6] 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 Powers ' trade blockade on
Chilean saltpeter.
History
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 [7] 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 chemists globally.
Haber, with his assistant Robert Le Rossignol , developed the high-pressure devices and catalysts needed to demonstrate the Haber process at laboratory scale. [8][9] 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. [5][10] 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. [5]
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.[11] During World War I , the production of
munitions required large amounts of nitrate. The Allies had access to large
sodium nitrate deposits in Chile (Chile saltpetre) controlled by British companies. Germany had no such resources, so the Haber process proved essential to the German war effort. [5][12] Synthetic ammonia from the Haber process was used for the production of nitric acid, a precursor to the nitrates used in explosives.
Process
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. [3]
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. [3]
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 (N 2) 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 principle .)
Kp (T ) for N 2 + 3 H 2 ⇌ 2 NH 3
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nickel catalyst . This reaction is used in the manufacture of ammonia gas and the process is called the Haber-Bosch process.
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