Yield of ammonia

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... 

Effect of Temperature and Pressure on the Yield of Ammonia in the Haber Process (PH2 = 3PNl)... 

Self-Test M.3 A In the synthesis of ammonia, what is the percentage yield of ammonia when 100. kg of hydrogen reacts with 800. kg of nitrogen to produce 400. kg of ammonia ... 

An important experimental task therefore remains to determine at least around the stoichiometrical composition, the variation of Kp with the composition. At high pressures the maximal yield of ammonia may occur at a hydrogen-nitrogen ratio different from three. 

C16-0019. Refer to Examples and. List four changes in conditions that might be used to increase the yield of ammonia in the Haber process. 

N2(g) + 3 H2(g) 2 NH3(g) Explain what effect an increase in pressure would have on the yield of ammonia. 

Lowering the temperature shifts the position of equilibrium to the right, so increasing the yield of ammonia, NH3. 

A temperature above 400 °C would increase the rate of the reaction, but would reduce the yield of ammonia. Below 400 °C, despite a higher yield, the rate of formation of ammonia is so slow as to be economically unviable. 

A pressure of 200 atmospheres increases the rate of attainment of equilibrium and also increases the yield of ammonia. The use of much higher pressures is prohibited by the high costs of building suitable containers that are able to withstand greater pressure in addition, the cost of the energy required to generate such large pressures is very expensive. 

The determination of nitrogen in the pure product was performed by the usual method of slow distillation with excess of dilute sodium hydroxide solution, the evolved ammonia being collected, during 4 hr. distillation, in standard acid, and determined by titration as usual, with methyl orange as indicator. The yield of ammonia corresponded to 18-19 per cent of nitrogen in this sample. Calculated for FCH2CONH, N= 18-18 per cent. The standard fiuoroacetamide was therefore 100 per cent pure. 

N2(g) + 3H2(g) 2NH3(g) AH° = —92 kj/mol Because the reaction is exothermic, heat is released as the reaction proceeds. Le Chatelier s principle predicts that the yield of ammonia is greater at lower temperatures. Just as in the contact process for manufacturing sulfuric acid, however, high yield is not the only important factor. The rate of reaction for ammonia synthesis is too slow at low temperatures. 

Le Chatelier s principle also predicts that the yield of ammonia is greater at higher pressures. High-pressure plants are expensive to huild and maintain, however. In fact, the first industrial plant that manufactured ammonia had its reaction vessel blow up. A German chemical engineer, Carl Bosch, solved this problem by designing a double-walled steel vessel that could operate at several hundred times atmospheric pressure. Modern plants operate at pressures in the range of 20 200 kPa to 30 400 kPa. 

By manipulating the pressure and removing ammonia from the reaction vessel, Haher successfully increased the yield of ammonia. To increase the rate of the reaction, Haher needed to find a catalyst. A catalyst would allow the reaction to proceed at higher temperatures—a compromise between rate and yield. Some historians claim that Haher performed more than 6500 experiments trying to find a suitable catalyst. He finally chose an iron catalyst. This catalyst works well at the relatively moderate temperature of 400°C that is used for the reaction. It lasts about five years before losing its effectiveness. 

For example, classic thermodynamic methods predict that the maximum equUi-brium yield of ammonia from nitrogen and hydrogen is obtained at low temperatures. Yet, under these optimum thermodynamic conditions, the rate of reaction is so slow that the process is not practical for industrial use. Thus, a smaller equilibrium yield at high temperature must be accepted to obtain a suitable reaction rate. However, although the thermodynamic calculations provide no assurance that an equUibrium yield will be obtained in a finite time, it was as a result of such calculations for the synthesis of ammonia that an intensive search was made for a catalyst that would allow equilibrium to be reached. 

It is implicit in reaction 9.4 that the equilibrium yield of ammonia is favored by high pressures and low temperatures (Table 9.1). However, compromises must be made, as the capital cost of high pressure equipment is high and the rate of reaction at low temperatures is slow, even when a catalyst is used. In practice, Haber plants are usually operated at 80 to 350 bars and at 400 to 540 °C, and several passes are made through the converter. The catalyst (Section 6.2) is typically finely divided iron (supplied as magnetite, Fe304 which is reduced by the H2) with a KOH promoter on a support of refractory metallic oxide. The upper temperature limit is set by the tendency of the catalyst to sinter above 540 °C. To increase the yield, the gases may be cooled as they approach equilibrium. 

Deamination appears to proceed equally rapidly both in the presence and absence of oxygen. Stein and Weiss (S18) examined this phenomenon and deduced that deamination can be accompanied by both oxidation and reduction. In the absence of oxygen the H atom, rather than the HO radical is the effective agent. This mechanism is satisfactory for the simple amino acids (glycine, alanine, and serine) studied by these workers. With the more complex amino acids yields of ammonia become less and side chains may be preferentially broken. 

The results of the calculations show that with increasing pressure the equilibrium yield of ammonia is increasing and that the non-ideality of the gas mixtures has in this case a positive effect on the equilibrium conversion. 

A gas-phase equilibrium responds to compression—a reduction in volume—of the reaction vessel. According to Le Chatelier s principle, the composition will tend to change in a way that minimizes the resulting increase in pressure. For instance, in the formation of NH3 from N2 and H2, 4 mol of reactant molecules produces 2 mol of product molecules. The forward reaction therefore decreases the number of molecules in the container and hence the pressure. It follows that, when the mixture is compressed, the equilibrium composition will tend to shift in favor of product, because that minimizes the increase in pressure (Fig. 9.12). Expansion results in the opposite response, a tendency for products to decompose. Haber realized that to increase the yield of ammonia, he needed to carry out the synthesis with highly compressed gases. The actual industrial process uses pressures of 250 atm and more (Fig. 9.13). 

We have seen how Haber s understanding of chemical equilibrium enabled him to increase the potential yield of ammonia (Fig. 9.15). He realized that he had to compress the gases and remove the ammonia as it was formed. As we have seen, compression shifts the equilibrium... 

For the system to move to a new state of equilibrium, Qc must increase, which means that the net reaction must go from left to right, as predicted by Le Chatelier s principle (Figure 13.11). In practice, the yield of ammonia in the... 

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