Ammonia synthesis equilibrium concentrations

Essential for synthesis considerations is the abiUty to determine the amount of ammonia present ia an equiUbrium mixture at various temperatures and pressures. ReHable data on equiUbrium mixtures for pressures ranging from 1,000 to 101,000 kPa (10 —1000 atm) were developed early on (6—8) and resulted ia the determination of the reaction equiUbrium constant (9). Experimental data iadicates that is dependent not only on temperature and pressure, but also upon the ratio of hydrogen and nitrogen present. Table 3 fists values for the ammonia equilibrium concentration calculated for a feed usiag a 3 1 hydrogen to nitrogen ratio and either 0 or 10% iaerts (10). 

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. 

Equilibrium. Forward and reverse reactions occurring at the same rate, resulting in a concentration of reactants. A + B C + D. Ammonia synthesis is an equilibrium reaction (N2 + 3H2 2NH3). 

Figure 17.20. Control of temperature in multibed reactors so as to utilize the high rates of reaction at high temperatures and the more favorable equilibrium conversion at lower temperatures, (a) Adiabatic and isothermal reaction lines on the equilibrium diagram for ammonia synthesis, (b) Oxidation of SOz in a four-bed reactor at essentially atmospheric pressure, (c) Methanol synthesis in a four bed reactor by the ICI process at 50 atm not to scale 35% methanol at 250°C, 8.2% at 300°C, equilibrium concentrations.

An experimental study of kinetics of ammonia synthesis on iron (101), cobalt, and nickel (96) catalysts, at ammonia concentrations much lower than that at equilibrium, showed that at pressures of the order of 1 atm the second of these possibilities is realized.7 When far from equilibrium, the... 

A mixture of reactants and products in the equilibrium state is called an equilibrium mixture. In this chapter, we ll address a number of important questions about the composition of equilibrium mixtures What is the relationship between the concentrations of reactants and products in an equilibrium mixture How can we determine equilibrium concentrations from initial concentrations What factors can be exploited to alter the composition of an equilibrium mixture This last question is particularly important when choosing conditions for the synthesis of industrial chemicals such as hydrogen, ammonia, and lime (CaO). 

In the early 1900s, the German chemist Fritz Haber discovered that a catalyst consisting of iron mixed with certain metal oxides causes the reaction to occur at a satisfactory rate at temperatures where the equilibrium concentration of NH3 is reasonably favorable. The yield of NH3 can be improved further by running the reaction at high pressures. Typical reaction conditions for the industrial synthesis of ammonia are 400-500°C and 130-300 atm. 

The law of mass action is widely applicable. It correctly describes the equilibrium behavior of all chemical reaction systems whether they occur in solution or in the gas phase. Although, as we will see later, corrections for nonideal behavior must be applied in certain cases, such as for concentrated aqueous solutions and for gases at high pressures, the law of mass action provides a remarkably accurate description of all types of chemical equilibria. For example, consider again the ammonia synthesis reaction. At 500°C the value of K for this reaction is 6.0 X 10 2 F2/mol2. Whenever N2, H2, and NH3 are mixed together at this temperature, the system will always come to an equilibrium position such that... 

For the ammonia synthesis reaction, the equilibrium expression can be written in terms of concentrations,... 

When the reactants and products of a given chemical reaction are mixed, it is useful to know whether the mixture is at equilibrium and, if it is not, in which direction the system will shift to reach equilibrium. If the concentration of one of the reactants or products is zero, the system will shift in the direction that produces the missing component. However, if all the initial concentrations are not zero, it is more difficult to determine the direction of the move toward equilibrium. To determine the shift in such cases, we use the reaction quotient (Q). The reaction quotient is obtained by applying the law of mass action, but using initial concentrations instead of equilibrium concentrations. For example, for the synthesis of ammonia,... 

It is important to understand the factors that control the position of a chemical equilibrium. For example, when a chemical is manufactured, the chemists and chemical engineers in charge of production want to choose conditions that favor the desired product as much as possible. In other words, they want the equilibrium to lie far to the right. When Fritz Haber was developing the process for the synthesis of ammonia, he did extensive studies on how the temperature and pressure affect the equilibrium concentration of ammonia. Some of his results are given in Table 6.2. Note that the amount of NH3 at equilibrium increases with an increase in pressure but decreases with an increase in temperature. Thus the amount of NH3 present at equilibrium is favored by conditions of low temperature and high pressure. 

To see how we can predict the effects of a change in concentration on a system at equilibrium, we will consider the ammonia synthesis reaction. Suppose there is an equilibrium position described by these concentrations ... 

Around 1900 Fritz Haber began to investigate the ammonia equilibrium [11] at atmospheric pressure and found minimal ammonia concentrations at around 1000 °C (0.012 %). Apart from Haber, Ostwald and Nernst were also closely involved in the ammonia synthesis problem, but a series of mistakes and misunderstandings occurred during the research. For example, Ostwald withdrew a patent application for an iron ammonia synthesis catalyst because of an erroneous experiment, while Nernst concluded that commercial ammonia synthesis was not feasible in view of the low conversion he found when he first measured the equilibrium at 50 - 70 bar [12] - [14],... 

Equations for describing ammonia synthesis under industrial operating conditions must represent the influence of the temperature, the pressure, the gas composition, and the equilibrium composition. Moreover, they must also take into consideration the dependence of the ammonia formation rate on the concentration of catalyst poisons and the influence of mass-transfer resistances, which are significant in industrial ammonia synthesis. 

Another test of validity is to check the performance of the model against experimental rate data obtained far from equilibrium. The microkinetic model presented in Table 7.3.1 predicts within a factor of 5 the turnover frequency of ammonia synthesis on magnesia-supported iron particles at 678 K and an ammonia concentration equal to 20 percent of the equilibrium value. This level of agreement is reasonable considering that the catalyst did not contain promoters and that the site density may have been overestimated. The model in Table 7.3.1 also predicts within a factor of 5 the rate of ammonia synthesis over an Fe(lll) single crystal at 20 bar and 748 K at ammonia concentrations less than 1.5 percent of the equilibrium value. 

A catalyst has no effect on the equilibrium composition. A catalyst increases the rates of both forward and reverse reactions to the same extent. The equilibrium composition and equilibrium concentration do not change when a catalyst is used, but the equilibrium composition is achieved in a shorter time. The role of a solid-phase catalyst in the synthesis of ammonia is shown in Figure 8.13. 

What does the equilibrium expression mean It means that, for a given reaction at a given temperature, the special ratio of the concentrations of the products to reactants defined by the equilibrium expression will always be equal to the same number—namely, the equilibrium constant K. For example, consider a series of experiments on the ammonia synthesis reaction... 

It is important to remember that although the changes we have just discussed may alter the equilibrium position, they do not alter the equilibrium constant. For example, the addition of a reactant shifts the equilibrium position to the right but has no effect on the value of the equilibrium constant the new equilibrium concentrations satisfy the original equilibrium constant. This was demonstrated earlier in this section for the addition of N2 to the ammonia synthesis reaction. 

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

Although the special ratio of products to reactants defined by the equilibrium expression is constant for a given reaction system at a given temperature, the equilibrium concentrations will not always be the same. Table 13.1 gives three sets of data for the synthesis of ammonia, showing that even though the individual sets of equilibrium concentrations are quite different for the different situations, the equilibrium constant, which depends on the ratio of the concentrations, remains the same (within experimental error). Note that subscript zeros indicate initial concentrations. 

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