Ammonia synthesis equilibrium position

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... 

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 ... 

These are the questions (among many others) to which thermodynamics gives an answer. However, the equilibrium positions predicted by thermodynamics may not always be attainable in practice. Indeed, in our example, the synthesis of ammonia, a catalyst is essential to facilitate the attainment of equilibrium. 

Let us consider the ammonia synthesis reaction. Suppose there is an equilibrium position described by these concentrations ... 

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. 

Consider again the ammonia synthesis reaction. The equilibrium constant always has the same value at a given temperature. At 500°C the value of is 6.0 X 10. Whenever N2, H2, and NH3 are mixed together at this temperature, the system will always come to an equilibrium position such that... 

In principle all reactions are reversible just as reactants have a tendency to combine and form products, products have the tendency to recombine and form the initial reactants. At equilibrium, the forward rate is balanced by the reverse rate, all net conversion ceases, and the composition of the system becomes constant in time. Suppose we load a closed reactor with a mixture that contains arbitrary amounts of the reactant and product species, and initiate the reaction while maintaining constant temperature and pressure. If we monitor the progress to equilibrium through the extent of reaction, we will observe it to increase in the positive or negative direction, indicating that the reaction progresses in the forward or reverse direction, until equilibrium is reached. Since temperature and pressure are held constant, the equilibrium state corresponds to conditions that minimize the Gibbs free ener. This condition allows us to obtain precise mathematical relationships for the equilibrium constant of the reaction. As an example, consider the ammonia synthesis reaction. 

The introduction of an inert gas, such as helium, into the reaction vessel for the synthesis of ammonia increases the total pressure in the vessel. But it does not change the partial pressures of the reaction gases present. Therefore, increasing pressure by adding a gas that is not a reactant or a product cannot affect the equilibrium position of the reaction system. 

Increase of reaction temperature on one hand accelerates the reaction speed but on the other hand, decreases the equilibrium ammonia concentration. The change of temperature affects the total rate of ammonia synthesis positively and negatively. Therefore there is an optimum reaction temperature at which the reaction velocity is the fastest and the synthesis efficiency is the highest. 

Haber and Le Rossignol " carried out further experiments to redetermine the equilibrium value using their same method at atmospheric pressure, but with various refinements. The principal reason why they carried out their measurements at atmospheric pressure was because they considered that their analytical methods were good enough. Furthermore, their experiment had the advantage of being able to approach the equilibrium state more easily from both sides, namely, synthesis and decomposition. They obtained a new value for the concentration of ammonia, 0.0048% at 1000 °C, which is close to the lower value (0.005%) of their former estimate. The reproducibility of their results was much better than that of Nemst. They were convinced that their results represented the true equilibrium positions more accurately than those of Nernst, since Nernst had not studied the decomposition of ammonia to complement the work on direct synthesis. 

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. 

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