In Industrial Use Ammonia is Continuously Removed
FIGURE 13.7 Represen- tation of the Haber process for the industrial production of ammonia. A mixture of gaseous N2 and H2 at 130-300 atm pressure is passed over a catalyst at 400-500°C, and ammonia is produced by the reaction N2(g) + 3 H2(g) 2NH3(y). The NH3 in the gaseous mixture of reactants and products is liquefied, and the unreacted N2 and H2 are recycled. [Pg.548]
They show catalytic activity (Chapter 7, p. 109) as elements and compounds. For example, iron is used in the industrial production of ammonia gas (Haber process, Chapter 11, p. 177). [Pg.155]
The industrial production of ammonia by use of natural gas feedstock can be represented by the following simplified set of reactions ... [Pg.93]
A process for the catalyzed industrial production of ammonia from N2 and H2 at high temperature and pressure. [Pg.23]
Change in Temperature Lack of Effect of a Catalyst Industrial Production of Ammonia... [Pg.540]
If, for instance, we add a catalyst to a mixture of PCI3 and CI2 at 523 K, the system will attain the same equilibrium concentrations of PCI3, CI2, and PCI5 more quickly than it did without the catalyst. Nevertheless, a catalyst often plays a key role in optimizing the yield of a reaction system. The industrial production of ammonia, described in the following subsection, provides an example of a catalyzed improvement of yield. [Pg.567]
This example highlights the distinction between thermodynamic and kinetic considerations. Even though NH3 forms spontaneously, it does so slowly in the industrial production of ammonia by the Haber process (Section 17.6), a catalyst is used to form NH3 at a practical rate. [Pg.664]
The industrial production of ammonia from natural gas involves eight different catalytic steps (Scheme 8-2). The overall reaction equation is given in Equation (8-1). [Pg.267]
The industrial production of ammonia is done by a process called Haber process. The reaction occurs as shown below, and is usually done in the presence of an iron catalyst. [Pg.157]
FIGURE 15.11 Diagram of the industrial production of ammonia. Incoming N2(g) and H2(g) are heated to approximately 500 °C and passed over a catalyst. When the resultant N2, H2, and NH3 mixture is cooled, the NH3 liquefies and is removed from the mixture, shifting the reaction to produce more NH3. [Pg.633]
Since the initial development of its synthesis, the industrial production of ammonia has been steadily increasing as can be seen from the curve (Figure 1), and in 2005 the world production has reached 147 million metric tonnes. Today 1.2% of the total world energy consumption is used for production of ammonia. [Pg.15]
In the Haber reaction, therefore, removing NH3 from an equilibrium mixture of N2, H2, and NH3 causes the reaction to shift right to form more NH3. If the NH3 can be removed continuously as it is produced, the yield can be increased dramatically. In the industrial production of ammonia, the NH3 is continuously removed by selectively liquefying it ( Figure 15.11). (The boiling point of NH3, —33 °C, is much higher than those of N2, —196 °C, and H2, —253 °C.) The liquid NH3 is removed, and the N2 and H2 are recycled to form more NH3. As a result of the product being continuously removed, the reaction is driven essentially to completion. [Pg.652]
The reaction Ibr the Haber process, the industrial production of ammonia, is... [Pg.467]
The industrial process of Haber—Bosch basically emulates the nitrogen fixation process to provide the resources for agricultural production at a faster pace to meet the rising demand. The industrial production of ammonia therefore provided a boost to production of nitrogen-based products. Ammonia is the basic building block for numerous chemical compounds downstream, such as conversion to polymers, fertilizers, etc. (Table 12.1). [Pg.278]
NH3 could be considered the compound of the twentieth century. On the one hand, it has been essential to the development of agriculture, but it also has played a major role in explosives (e.g., TNT). Just as it has been essential to the welfare of billions of people, it is also estimated to have caused the deaths of around 150-200 million people. There is a long list of processes and equipment that have been made possible by the industrial production of ammonia, such as, for example, nylon, cattle feed, cleaning equipment, and air conditioning. [Pg.52]
In 1912 the German chemist Fritz Haber (1868-1934) developed a process for S5nthesizing ammonia directly from nitrogen and hydrogen (Figure 15.4 ). The process is sometimes called the Haber-Bosch process to also honor Karl Bosch, the engineer who developed the equipment for the industrial production of ammonia. The engineering needed to implement the Haber process requires the use of temperatures and pressures (approximately 500°C and 200 atm) that were difficult to achieve at that time. [Pg.579]
Even though catalysts cannot change the reaction yield, they often play key roles in optimizing it. The industrial production of ammonia, described in the next subsection, provides an example of a catalyzed improvement of yield. [Pg.568]
| Figure 18.13 Schematic diagram of the process for the industrial production of ammonia from nitrogen and hydrogen. |  |
A study of the reaction N, + 3H, - 2NH" used for the industrial production of ammonia, reveals that the rate of this reaction is actually controlled by the rate at which nitrogen dissociates into atomic nitrogen. Thus, anything that increases the rate ofthe reaction N, - 2N automatically will increase the production rate of ammonia. This also reveals the most important point in the study of reaction mechanisms When a reaction proceeds through more than one step, the slowest one determines the rate ofthe overall reaction and therefore determines the observed rate equation. The slowest step, the bottleneck in the series, is said to be rate determining. These ideas may be illustrated by the following analogy ... [Pg.414]
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