The production of ammonia from its elements

The fixing of nitrogen in ammonia is the first step in the industrial production of mairy nitrogen-containing materials, such as fertilizers. The industrial 

The indusU ial reaction using nitrogen and hydrogen gases 

The reverse reaction to ammonia synthesis, the decomposition to nitrogen and hydrogen, is used in die nitriding of iron and canied out industiially at temperatures around 800 K and atmospheric pressure to produce surfacehardening. This dissolution reaction must also play a part in the synthesis of ammonia by the industiial process. The attempt to ninide non by reaction with nin ogen gas is vety slow under atmospheric pressure, presumably due to the stability of the nitrogen molecule. 

The kinetic parameters for NH3 decomposition at high pressures and temperature around 650 K are found to be 

The overall rate for tire formation of ammonia must tlrerefore be a balance between the formation and tire decomposition of the product species. Experimental data suggest tlrat tlris balance can be represented by the equation 

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Under prevailing conditions it was found that x = 0.5275, whence Kp = 6.79 X 10 2 bar . Thus, under these conditions the production of ammonia from its elements is far from complete. 

It will be seen from the above equation that large amounts of energy are mot required for the production of ammonia from its elements, and that, therefore, the manufacture of synthetic ammonia need not be confined to districts where large amounts of cheap water-power are available, as is the case with the electrical production of nitric acid, cyanamide, etc. 

The synthesis of ammonia from its elements, N2 + 3H2 2NH3, is one of the largest industrial processes based on heterogeneous catalysis [1-4]. Almost 90% of the present world capacity of about 200 million tons per year is used for the production of fertilizers. Almost all plants employ promoted catalysts based on iron, quite similar to the one developed by A. Mittaschin 1910 attheBadische Anilin and Sodafabrik (BASF) [5]. The above reaction could be successfully realized for the first time in the laboratory in 1909 by Fritz Haber, and was then transferred into a technical... 

In 1918, Fritz Haber won the Nobel Prize for the synthesis of ammonia from its elements. In other words, Fritz Haber was the first chemist to create ammonia in a laboratory. His discovery made it possible to manufacture industrial-sized quantities of ammonia that could be added to fertilizers. (Check out Chapter 17 for details on why ammonia is a useful fertilizer.) Manufacturing ammonia led to increased crop yields and transformed farming, shaping the production of foods you eat today. It s arguably the most important development in agriculture since the development of irrigation (6,000 BC) and the use of draft animals (4,000 BC). 

Fritz Haber, 1968-1934, German physical chemist, winner of the Nobel Prize for chemistry in 1918 for the synthesis of ammonia from its elements. With Carl Bosch, he invented a process for the large scale production of ammonia for nitrogen fertilizers. 

Process for synthetic production of ammonia from its elements, whereby an appropriate mixture of nitrogen and hydrogen is continuously subjected to both the production of ammonia under the influence of heated catalysts, and continuous removal of the resulting ammonia, characterized by constant pressure and the transfer of process heat from the ammonia-containing reaction gases to the incoming ammonia-free gas mixture. ... 

Haber s third key patent for ammonia synthesis—No. 238450 filed on September 14, 1909, and issued on September 28, 1911, in Haber s, not BASF s, name, and generally known as the high-pressure patent—stressed the catalysis under pressure and high temperature. Its claim read Process for production of ammonia from its elements using catalysis under pressure and elevated temperature, whose characteristic is that the synthesis takes place under very high pressure of about 100 atm, but to be useful at 150-250 atm or higher. 

The production of ammonia is of historical interest because it represents the first important application of thermodynamics to an industrial process. Considering the synthesis reaction of ammonia from its elements, the calculated reaction heat (AH) and free energy change (AG) at room temperature are approximately -46 and -16.5 KJ/mol, respectively. Although the calculated equilibrium constant = 3.6 X 108 at room temperature is substantially high, no reaction occurs under these conditions, and the rate is practically zero. The ammonia synthesis reaction could be represented as follows ... 

Haber demonstrated that the production of ammonia from the elements was feasible in the laboratory, but it was up to Carl Bosch, a chemist and engineer at BASF, to transform the process into large-scale production. The industrial converter that Bosch and his coworkers created was completely revised, including a cheaper and more effective catalyst based on extensive studies in high-pressure catalytic reactions. This approach led to Bosch receiving the Nobel Prize in chemistry in 1931, and the production of multimillion tons of fertilizer per year worldwide, see also Agricultural Chemistry Catalysis and Catalysts Equilibrium Le Chatelier, Henri Nernst, Walther Hermann Ostwald, Friedrich Wilhelm. 

Another critical use of nitrogen is in the production of ammonia by the Haber process, named after its inventor, the German chemist Fritz Haber. The Haber process involves the direct synthesis of ammonia from its elements, nitrogen and hydrogen. The two gases are combined at temperatures of 932-1,292°F (500-700°C) under a pressure of several hundred atmospheres over a catalyst such as finely divided nickel. One of the major uses of the ammonia produced by this method is in the production of synthetic fertilizers. 

Although the direct synthesis of ammonia from its elements had been known for some time, the yield of product was found to be negligible. In H 1905, Fritz Haber (1868-1934) began to study this reaction, employing the... 

Analyze and Plan We are asked to calculate the entropy change for the synthesis of ammonia from its constituent elements. Standard molar entropy values for the reactants and the product are given in Table 19.2. 

One of the interesting and useful features of chemical equilibria is that they can be manipulated in specific ways to maximize production of a desired substance. Consider, for example, the industrial production of ammonia from its constituent elements by the Haber process ... 

At that time, the possibility of producing chemically even traces of ammonia from its elements was disputed. Only the method of electrical production by silent discharge was proven.. . . Of course, the stability of the compound at ordinary temperatures seemed probable, in view of the amount of heat necessary for its formation. But after the numerous fmitless experiments of the previous generation, it seemed beyond doubt that nitrogen and hydrogen would not unite spontaneously at this temperature. Experiments at high temperatures, on the other hand, had up to that time proved nothing but its rate of decomposition, and... 

The previous sections mainly considered the individual process steps involved in the production of ammonia and the progress made in recent years. The way in which these process components are combined with respect to mass and energy flow has a major influence on efficiency and reliability. Apart from the feedstock, many of the differences between various commercial ammonia processes lie in the way in which the process elements are integrated. Formerly the term ammonia technology referred mostly to ammonia synthesis technology (catalyst, converters, and synthesis loop), whereas today it is interpreted as the complete series of industrial operations leading from the primary feedstock to the final product ammonia. 

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