Haber process equilibrium

Although the left to right reaction is exothermic, hence giving a better equilibrium yield of sulphur trioxide at low temperatures, the reaction is carried out industrially at about 670-720 K. Furthermore, a better yield would be obtained at high pressure, but extra cost of plant does not apparently justify this. Thus the conditions are based on economic rather than theoretical grounds (cf Haber process). 

The Haber process, represented by this equation, is now the main source of fixed nitrogen. Its feasibility depends on choosing conditions under which nitrogen and hydrogen react rapidly to give a high yield of ammonia. At 25°C and atmospheric pressure, the position of the equilibrium favors the formation of NH3 (K= 6 x 105). Unfortunately. however, the rate of reaction is virtually zero. Equilibrium is reached more rapidly by raising the temperature. However, because... 

As an indispensable source of fertilizer, the Haber process is one of the most important reactions in industrial chemistry. Nevertheless, even under optimal conditions the yield of the ammonia synthesis in industrial reactors is only about 13%. This Is because the Haber process does not go to completion the net rate of producing ammonia reaches zero when substantial amounts of N2 and H2 are still present. At balance, the concentrations no longer change even though some of each starting material is still present. This balance point represents dynamic chemical equilibrium. 

In this chapter, we present basic features of chemical equilibrium. We explain why reactions such as the Haber process cannot go to completion. We also show why using catalysts and elevated temperatures can accelerate the rate of this reaction but cannot shift Its equilibrium position in favor of ammonia and why elevated temperature shifts the equilibrium In the wrong direction. In Chapters 17 and 18, we turn our attention specifically to applications of equilibria. Including acid-base chemistry. 

What happens to a system at equilibrium if the concentration of one of the reacting chemicals is changed This question has practical importance, because many manufacturing processes are continual. Products are removed and more reactants are added without stopping the process. For example, consider the Haber process that was mentioned in the previous Sample Problem. 

In this chapter, you learned about the Haber process for manufacturing ammonia. You used this process to help you understand various concepts related to equilibrium. As you can see in Figure 7.11, ammonia is a valuable industrial chemical. Its annual global production is well over 100 million tonnes. The vast majority of ammonia, roughly 80%, is used to make fertilizers. You will now examine how the equilibrium concepts you have been studying work together to provide society with a reliable, cost-effective supply of ammonia. 

According to Le Chatelier s Principle, the production of ammonia is favored by a high pressure and a low temperature. The Haber process is typically carried out at pressures between 200 and 400 atmospheres and temperatures of 500°C. While Le Chatelier s Principle makes it clear why a high pressure would be favorable in the Haber process, it is unclear why a high temperature would be desirable because the reaction is exothermic. An increase in temperature shifts an exothermic reaction to the left. Even though the equilibrium shifts to... 

One of the principal goals of chemical synthesis is to maximize the conversion of reactants to products while minimizing the expenditure of energy. This objective is achieved easily if the reaction goes nearly to completion at mild temperature and pressure. If the reaction gives an equilibrium mixture that is rich in reactants and poor in products, however, then the experimental conditions must be adjusted. For example, in the Haber process for the synthesis of ammonia from N2 and H2 (Figure 13.7), the choice of experimental conditions is of real economic importance. Annual U.S. production of ammonia is about 13 million tons, primarily for use as fertilizer. 

Let s consider the equilibrium that occurs in the Haber process for the synthesis of ammonia ... 

Figure 15.4 Equilibrium mol % NH3 produced in the Haber process as a function of temperature at several different pressures, as reported by Haber in his Nobel Prize address (See F. Haber, Naturwissenschaften, 49, 1041-1049 (1922).)...

The manner in which this equilibrium is influenced by changes in temperature and pressure should be reviewed. In practice, ammonia is produced by the Haber process at temperatures ranging from 400 to 600°C and at pressures between 200 and 1000 atm. Catalysts that are suitable for use in this process include a mixture of the oxides of iron, potassium, and aluminum iron oxide alone mixtures of iron and molybdenum the metals platinum, osmium, uranium and a number of others as well. 

Example 6. Write down the form of the equilibrium constant that applies to the Haber process for the production of ammonia. 

One of the most important industrial processes is the Haber process, based on the reaction N2 + 3H2 -e- 2NH3. Calculate the equilibrium constant of this reaction at 298 K and, using temperature-independent heat capacities, at 1000K. In light of your result, can you guess why the reaction is carried out at a high temperature ... 

Le Chitelier s principle states that if an equilibrium system is stressed, the equilibrium will shift in the direction that relieves the stress. For the Haber process at equilibrium, if more nitrogen is added, then the process will shift to the right. If nitrogen is removed, then the equilibrium will shift to the left. 

The following equilibrium concentrations were observed for the Haber process at 127°C ... 

Obviously, the kinetics and the thermodynamics of this reaction are in opposition. A compromise must be reached, involving high pressure to force the equilibrium to the right and high temperature to produce a reasonable rate. The Haber process for manufacturing ammonia represents such a compromise (see Fig. 19.6). The process is carried out at a pressure of about 250 atm and a temperature of approximately 400°C. Even higher temperatures would be required if a catalyst consisting of a solid iron oxide mixed with small amounts of potassium oxide and aluminum oxide were not used to facilitate the reaction. 

Unfortunately, although a low temperature increases the yield of ammonia, the reaction becomes uneconomically slow, as now there are fewer sufficiently energetic collisions between the particles. Obviously the industrial chemist faces the problem of reconciling both the kinetic and equilibrium considerations for this reaction to make economical to operate. In practice a compromise is reached by selecting an optimum operating temperature and pressure (see the Haber Process on page 248). 

The production of ammonia by the Haber Process utilises all four factors that affect equilibrium reactions ... 

The chemical industry, such as the factory in Figure 15, makes use of Le Chatelier s principle in many ways. One way is in the synthesis of ammonia by the Haber Process. High pressure is used to drive the following equilibrium to the right. 

In the Haber Process, how would removing NH3 as it forms affect the equilibrium ... 

Although the equilibrium lies far to the right (K 106), a large amount of energy ( a) is required to break the N2 triple bond. Because of the exothermicity, however, K decreases significantly when the temperature is increased. The Haber process requires the use of an iron-based catalyst at 250 atm and 400 °C. 

Factors That Affect Equilibria 17-7 The Haber Process A Practical Application of Equilibrium 17-8 Application of Stress to a System at Equilibrium... 

Nitrogen, N2, is very unreactive. The Haber process is the economically important industrial process by which atmospheric N2 is converted to ammonia, NH3, a soluble, reactive compound. Innumerable dyes, plastics, explosives, fertilizers, and synthetic fibers are made from ammonia. The Haber process provides insight into kinetic and thermodynamic factors that influence reaction rates and the positions of equilibria. In this process the reaction between N2 and H2 to produce NH3 is never allowed to reach equilibrium, but moves toward it. 

Knowing the factors that affect chemical equilibrium has great practical value for industrial applications, such as the synthesis of ammonia. The Haber process for synthesizing ammonia from molecular hydrogen and nitrogen uses a heterogeneous catalyst to speed up the reaction (see p. 540). Let us look at the equilibrium reaction for ammonia synthesis to determine whether there are factors that could be manipulated to enhance the yield. 

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