Haber Process

10 Key excerpts on "Haber Process"

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  eBook - ePub

  Chemistry II For Dummies

  • John T. Moore(Author)
  • 2012(Publication Date)
  • For Dummies

    (Publisher)

  The Haber-Bosch Process, often just referred to as the Haber Process, is basically the process by which ammonia gas is produced from hydrogen and nitrogen gases. This chemical process is one of the most important in the world. It is the first step is in the production of synthetic fertilizers. Without this process, famine would spread across the globe.

  In the modern Haber-Bosch process, atmospheric nitrogen is produced from liquefying air and the hydrogen gas is produced by reacting methane with steam: The nitrogen and hydrogen gases combine exothermically to produce ­ammonia:

  The equilibrium constant expression, Kc , for this reaction at 25°C is 4 × 108 .

  This large value of the equilibrium constant indicates that a large quantity of ammonia should be produced, but this reaction is so slow that not much ammonia is produced in a reasonable period of time. If the temperature is increased in order to increase the speed of reaction (kinetics), the Kc decreases dramatically. In order for the reaction to be economically practical, LeChatelier’s Principle must be applied.

  LeChatelier’s Principle allows chemists to adjust the concentrations of the reactants, the pressure, and the temperature in order to increase the amount of ammonia produced. This increased efficiency allows more fertilizer to be made at a lower price. This is certainly important to US farmers but is of critical importance to farmers in third world countries.

  The history of the Haber Process

  In 1909, a German chemist, Fritz Haber, developed a method of combining nitrogen gas with hydrogen gas under easily achievable conditions to produce ammonia, NH3

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  Hydrogenation

  Catalysts and Processes

  • S. David Jackson(Author)
  • 2018(Publication Date)
  • De Gruyter

    (Publisher)

  Justin S. J. Hargreaves

  7 Heterogeneously catalyzed ammonia synthesis

  7.1 Introduction

  The development of the Haber–Bosch process was undoubtedly a landmark of the twentieth century. Through the provision of access to synthetic fertilizers, this single process has been directly credited with sustenance of a significant fraction of the global population. Indeed, nitrogen fixation from natural means accounts for about 50% of that necessary for plant and crop growth with the Haber–Bosch process accounting for the remainder [1 , 2 ]. This is a fact which is reflected in the nitrogen content within the bodies of each and everyone of us in that a significant proportion will have passed through Haber–Bosch process. Further statistics associated with the process are that it is estimated, when considered in its entirety, to be responsible for the consumption of 1–2% of manmade energy [3 ] and to be responsible for a significant proportion of the global manmade carbon dioxide emissions. In terms of the latter point, the CO2 intensive nature of the process arises in the use of fossil fuel sources to derive hydrogen and in their application as energy source. In 2010, the process was operated at the scale of 131 million tons of ammonia production, releasing an associated 245 million tons of CO2 [4 ].

  In general industrial operation, the process involves the reaction between nitrogen and hydrogen over a promoted iron-based catalyst. The feedstreams have to be of high purity as the catalyst is very susceptible to poisoning by even very small traces of oxygenates. As can be seen below, the reaction is favored by high pressure, and reaction pressures of >100 atmospheres are applied on the industrial scale.

  1 / 2

  N 2

  + 3 / 2

  H 2

  ↔

  NH 3

    Δ

  H °

  = − 46 kJ .

  mol -1

  Although the reaction is thermodynamically favored at low reaction temperatures, ca. 400–500°C is applied in practice to achieve acceptable process kinetics. Table 7.1 , taken from Fritz Haber’s Nobel Prize lecture [5

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  Regulation, Market Prices, And Process Innovation

  The Case Of The Ammonia Industry

  • Edward Greenberg, Christopher T. Hill, David J. Newburger(Authors)
  • 2019(Publication Date)
  • Taylor & Francis

    (Publisher)

  *

  4.2 Ammonia Synthesis

  Today virtually all ammonia is produced commercially by direct synthesis from hydrogen and nitrogen by means of the Haber-Bosch process. The very simple chemistry of the process can be described by the chemical equation of equilibrium:

  N2 + 3H2 ⇄ 2NH3

  For reasons discussed below, the reaction is carried out in the presence of a catalyst at high temperature and pressure using a gas mixture which contains relatively small amounts of certain impurities.

  The optimum operating temperature, pressure, and catalyst for ammonia synthesis are dependent upon two major concerns: the equilibrium conversion of N2 and H2 to NH3 , and the rate at which that conversion is reached. The chemical relationship between the three gases is an equilibrium one, which means that after a sufficient time has elapsed, each of the gases will be present at a fractional concentration that depends only on pressure and temperature. Haber and others made careful measurements of this equilibrium and found that the concentration of NH3 is enhanced by low temperature and high pressure, as shown in Table 4.1 taken from Slack and James. [1]

  That this should be the case follows from two facts: first, assuming ideal gas behavior, the reaction of one volume of N2 and three volumes of H2 produces only two volumes of NH3

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  Chemical Thermodynamics: Advanced Applications

  Advanced Applications

  • J. Bevan Ott, Juliana Boerio-Goates(Authors)
  • 2000(Publication Date)
  • Academic Press

    (Publisher)

  Maybe this solution is not final. The nitrogen bacteria teach [us] that nature in the refined forms of the chemistry of life still knows and realizes possibilities, the imitation of which is at present beyond our ability. It suffices that in the meantime, a new abundance of food benefits mankind from a richer nitrogen fertilization of the soil, and that the chemical industry is coming to the aid of the farmer who changes stones into bread on this peaceful earth

  l

  Today ammonia is routinely synthesized by the Haber Process. Pressures of 300 to 500 bar (or atm) and temperatures of 700 to 800 K are used. The usual catalyst contains about 95 wt% Fe3 O4 , along with some K2 O (∼ 1 wt%), Al2 O3 (∼2 wt%), CaO (∼2 wt%),m and smaller amounts of MgO and SiO2 . Containment is achieved using a steel reaction vessel to give strength, but lined with soft iron to prevent embrittlement of the steel due to contact with the H2 .n The procedure followed can be represented by the flow chart shown in Figure 15.5 . A mixture of N2 (g) and H2 (g) is compressed, then put in the reaction vessel where it is heated over the catalyst. Reaction occurs to produce an equilibrium amount of NH3 . This equilibrium mixture then leaves the reaction vessel (and catalyst) and is cooled so that the ammonia condenses to liquid and is removed. The remaining N2 (g) and H2 (g) is then returned to the reaction mixture. The final result is an essentially complete conversion of the (H2 + N2 ) gaseous mixture into NH3 .

  Figure 15.5 The Haber cycle used to produce liquid ammonia.

  The importance of this reaction can be understood by recognizing that over 20 million tons of NH3 were produced in the United States by this process in 1998. This amount can be compared with the production of sulfuric acid (48 million tons), the chemical that is produced in the largest amount.o

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  Six Chemicals That Changed Agriculture

  • Robert L Zimdahl(Author)
  • 2015(Publication Date)
  • Academic Press

    (Publisher)

  Between 1886 and 1891, he studied at the University of Heidelberg, the University of Berlin, and the Technical College of Charlottenburg, Germany. After his early studies, he worked in his father’s chemical business and in the Swiss Federal Institute of Technology in Zürich. After a few years, he moved to the Technische Hochschule’s (Institute of Technology) Department of Chemical and Fuel Technology at the University of Karlsruhe, Germany (founded in 1825) where he worked as a physical chemist from 1894 to 1911. In 1905, Haber accomplished successfully what chemists had long sought to do—fix N 2 from air and combine it with hydrogen (H 2) gas to synthesize ammonia (NH 3). He used high pressure and a catalyst (iron + potassium hydroxide [KOH]). Carl Bosch, 5 a chemist and engineer employed by Badische Anilin und Soda-Fabrik (BASF) studied high pressure chemistry between 1909 and 1913. In 1910, he succeeded in transforming Haber’s tabletop method of fixing nitrogen into an important industrial process that eventually produced megatons of fertilizer and much smaller amounts of explosives. The Haber–Bosch process was separated production of nitrogen products (e.g., fertilizer, explosives, chemical feedstocks) from natural nitrate deposits. The Haber–Bosch Process The Haber–Bosch process is the catalytic synthesis of ammonia from the hydrogen in natural gas (CH 4) and inert atmospheric nitrogen under high temperature (400–450 °C) and pressure (200 atm = approximately 3000 psi). Natural gas is 70–90% of the cost of the process. For several years the United States imported natural gas primarily from Trinidad (45%) and Canada (23%). With recent increased production of natural gas, imports have declined and less than 3% of natural gas that supplies 24% of US energy is now imported

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  Encyclopedic Dictionary of Named Processes in Chemical Technology

  • Alan E. Comyns(Author)
  • 2014(Publication Date)
  • CRC Press

    (Publisher)

  147 H H H See Pechiney H. Haber Also called Haber–Bosch and Haber–Bosch–Mittasch . A process for synthesizing ammonia from the elements, using high temperatures and pressures and an iron-containing catalyst. Invented by F. Haber at BASF in 1908. In 1909, C. Bosch of BASF built a pilot plant using an osmium-based catalyst; in 1913, a larger plant was built at Oppau, Germany. The process has been continually improved and is still of major importance worldwide. Haber was awarded the Nobel Prize for this work in 1918 but was infamous for his introduction of poison gases in World War I. Bosch was awarded the Nobel Prize in 1931. German Patents 235,421; 293,787. Haber, F. and van Oordt, G., Z. Anorg. Allg. Chem ., 1905, 43 , 111. Harding, A.J., Ammonia: Manufacture and Uses , Oxford University Press, London, U.K., 1959. Chilton, T.H., Strong Water: Nitric Acid, Its Sources, Methods of Manufacture, and Uses , MIT Press, Cambridge, MA, 1968, 64. Haber, L.F., The Chemical Industry 1900–1930 , Clarendon Press, Oxford, U.K., 1971, 187. Vancini, C.A., Synthesis of Ammonia , translated by L. Pirt, Macmillan Press, Basingstoke, U.K., 1971, 234. Jennings, J.R. and Ward, S.A., in Catalyst Handbook , 2nd ed., Twigg, M.V., Ed., Wolfe Publishing, London, U.K., 1989, 384. Travis, T., Chem. Ind. (London) , 1993, (15), 581. Leigh, G.J., The World’s Greatest Fix: A History of Nitrogen and Agriculture , Oxford University Press, Oxford, U.K., 2004, 129–134. Stoltzenerg, D., Fritz Haber: Chemist, Nobel Laureate, German, Jew , Chemical Heritage Press, Philadelphia, PA, 2004, Chap. 5. Haber Gold A hydrometallurgical process for extracting gold from its ores. It uses a proprietary extracting agent that is specific for gold, instead of cyanide. Developed by Haber Inc. and proposed for installation by Gold City Inc. at its Winnemucca, NV, plant in 2004. Chem. Eng. (N.Y.) , 2004, 111 (2), 14.

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  Catalyst Handbook

  • Martyn V. Twigg(Author)
  • 2018(Publication Date)
  • Routledge

    (Publisher)

  Figure 8.1 which traces the energy required for the fixation of nitrogen during this century. The introduction of the Haber Process, using coal, gave rise to a substantial improvement in the energy requirement. Subsequently the improvement obtained was small as the coal based process was refined, and it was not until the introduction of natural gas and naphtha based plants followed by the large single stream plants that further significant reductions in energy consumption were achieved. More recently the development of the high efficiency processes has brought the energy consumption close to the theoretical minimum for processes based on hydrocarbon feedstocks.

  Figure 8.1 Improved efficiency of nitrogen fixation during this century.

  The development of the science required for ammonia synthesis started in the 19th century, and its subsequent rapid progress owes much to Haber,

  232 ,233

  who first determined the equilibrium constant for the synthesis reaction in 1904. He conducted his experiments at 1020°C and atmospheric pressure over finely divided iron catalyst, and he was able to extrapolate the data obtained to lower temperatures using newly developed theories in physical chemistry. He concluded that it would not be possible to develop a viable process because conversions were so low. In 1906 Nernst234 applied his new heat theory, which we now know as the Third Law of Thermodynamics, to the synthesis reaction. From heats of reaction he was able to calculate the equilibrium constant at any temperature and pressure. Moreover, he also determined a value at 75 bar, which appeared to substantiate the value he had obtained from his theory and to refute the earlier data obtained by Haber. In the ensuing debate Haber and Le Rossignol235 undertook further experimental work, this time at pressures up to 30 bars, and showed that the equilibrium constant was higher that that predicted by Nernst. He attributed this to the uncertainty of the value used by Nernst for the specific heat of ammonia at high temperatures. More importantly, Haber decided that even though conversions were low, a process to synthesize ammonia was feasible provided that: (1) synthesis and removal of the ammonia formed were carried out at high pressure; (2) synthesis gas was recirculated over the catalyst and (3) a satisfactory catalyst could be developed. Haber and his co-workers developed a catalyst based on osmium, and in 1909 built a semi-technical plant capable of producing 90 g h−1 of ammonia. Following the successful demonstration of the unit, BASF decided to undertake large-scale development work and, with Carl Bosch as project leader, a plant was built at Oppau to produce 30 tonnes day−1

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  From Stars To Stalagmites: How Everything Connects

  • Paul S Braterman(Author)
  • 2012(Publication Date)
  • World Scientific

    (Publisher)

  Throughout his career, Haber supplemented his income by industrial consulting, and his work with BASF was to lead to the achievement for which he will always be remembered. Simple Reaction, Complex Problem, Major Importance At this stage, as an expert on thermodynamics, and on the industrial reac-tions of gases, Haber was naturally drawn to the outstanding chemical problem of the decade: the formation of ammonia directly from its ele-ments nitrogen and hydrogen. This means taking one molecule of nitrogen (N 2 ) and three of hydrogen (H 2 ), and reacting them together to make two molecules of ammonia (NH 3 ). This ammonia can be converted to nitric acid and nitrates, using processes already well established at that time. 90 From Stars to Stalagmites: How Everything Connects The principal use of nitrate was as a fertiliser. This had been known for centuries, and nitrate use was fundamental to the agricultural revolu-tion of the mid-19th century. In 1900 the main source of nitrates and nitric acid was Chile saltpetre, naturally occurring impure sodium nitrate derived ultimately from bird and bat excrement. This material, which had formed vast accumulations under the extremely dry conditions of the Atacama Desert, provided 60% of all nitrogenous fertiliser worldwide, and Germany was importing one third of all Chilean production. The Chilean deposits are extensive, but not unlimited, and their essen-tial role in agriculture was already foreseen as a serious long-term prob-lem. In 1898, William Crookes, President of the British Association for the Advancement of Science, had warned that some way must be found to convert atmospheric nitrogen to nitrate, saying “It is the chemist who must come to the rescue... It is through the laboratory that starvation may ultimately be turned to plenty.” But that was not Haber’s main motivation.

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  The Chemists' War

  1914-1918

  • Michael Freemantle(Author)
  • 2015(Publication Date)
  • Royal Society of Chemistry

    (Publisher)

  th century, not just in Germany but in the whole world, and for several reasons. He is most widely known for the process that is named after him. Well before the First World War, he developed a workable process for converting chemically-free nitrogen in the atmosphere into chemically-fixed nitrogen in the form of ammonia. German industrial chemist Carl Bosch (1874–1940) and colleagues subsequently developed his nitrogen fixation process and used the ammonia to manufacture nitrogenous fertilizers in vast quantities for use in agriculture.

  Haber had effectively found a way of making “bread from air.” The Haber–Bosch ammonia synthesis process, as it became known, helped feed millions of people around the world and, according to some commentators, was largely responsible for averting famine and saving the world from starvation.

  3 ,4

  It detonated the population explosion in the 20th century, observed Vaclav Smil, a Czech-Canadian scientist and policy analyst at the University of Manitoba, Canada. Writing in 1999, Smil suggested that “the world’s population could not have grown from 1.6 billion in 1900 to today’s six billion,” if it had not been for the process.5 Without the process, “almost two-fifths of the world’s population would not be here” he explained.

  The ammonia synthesis process also played a critical role in the war. In the same way that Britain would have found it difficult to sustain its war effort for four years without the acetone produced by the Weizmann process (see Chapter 7 ), Germany could not have continued fighting the war beyond a few months without the Haber–Bosch process that enabled it to continue manufacturing nitrogen-containing explosives such as trinitrotoluene, nitroglycerine, and nitrocellulose (see Chapter 9 ). The process therefore facilitated the production of not just “bread from air” but also “explosives from air.”

  Haber’s nitrogen fixation process was not his only contribution to the applied chemistry of the 20th century. He is also well known and much vilified for spearheading Germany’s development of lethal chemical weapons and overseeing their introduction on the Western Front in April 1915 and their continued use throughout the war. The “Father of Modern Chemical Warfare,” as Haber became known, considered chemical warfare to be humane (see Chapter 13 ). He was not alone in this view. Munitions and war experts in both the Central Powers and the Allied Powers shared this “perverse-sounding notion” notes Bretislav Friedrich, physics professor at the Fritz Haber Institute of the Max Planck Society in Berlin.6

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  Periodic Operation of Chemical Reactors

  • P. L. Silveston, R. R. Hudgins(Authors)
  • 2012(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  Chapter 2 Hydrogenation Processes

  Robert Ross Hudgins and Peter Lewis Silveston, Waterloo, Ontario, Canada

  Outline

  2.1 Ammonia Synthesis 2.1.1 Iron Catalysts

  2.1.1.1  Mass Transfer Interference

  2.1.1.2  Near-Adiabatic Operation

  2.1.2 Ruthenium Catalyst 2.1.3 Osmium Catalyst 2.1.4 Interpretation

  2.1.4.1  Relaxed Steady State

  2.1.4.2  Reactant Inhibition

  2.1.4.3  Reactant Storage

  2.1.4.4  Surface Activation/Restructuring

  2.1.5 Comments on the Ammonia Synthesis Application

  2.2 NOx Reduction

  2.3 Methanation 2.4 Methanol Synthesis 2.4.1 Copper-Zinc-Alumina Catalysts 2.4.2 Copper Zinc Catalyst 2.4.3 Commercial Methanol Catalysts 2.5 Ethylene Hydrogenation 2.6 Aromatics Hydrogenation 2.7 Oscillatory Behavior

  In this chapter, the hydrogenation of small molecules, such as nitrogen, nitrogen oxides, carbon monoxide, ethylene and the simple aromatics are examined under periodic operation. under periodic operation are examined. There is no mention in the literature of the application of modulation to the hydrogenation of complex molecules such as found in hydrocracking, hydrodesulfurization, hydrodenitrogenation or in the hydrotreating of polymers, though they certainly warrant attention. The Fischer-Tropsch synthesis may be considered a hydrogenation reaction, but it is also a polymerization one, so it will be discussed in Chapter 7 .

  2.1 Ammonia Synthesis

  In industrial hydrogenation processes, ammonia synthesis is among the most important. Between 1980 and 2000, some 15 studies of the periodic operation of this synthesis were published. Employing a triply promoted iron catalyst, it is one of the largest commercial applications of catalysis.

  Thermodynamics suggest that low temperatures should be used, but rates are far too slow for practical purposes. At higher temperatures where rates are much higher, conversion is equilibrium-limited, usually to less than 5%. This limitation is overcome by taking advantage of the stoichiometry and conducting the synthesis at pressures as high as 10 MPa, but high pressure and low conversion-per-pass makes reactor operation expensive.

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