Haber process reaction rate

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. 

In reactions involving only gases, for example the Haber process (Chapter 11, p. 177), an increase in the overall pressure at which the reaction is carried out increases the rate of the reaction. The increase in pressure results in the gas particles being pushed closer together. This means that they collide more often and so react faster. 

In a catalytic experiment involving the Haber process, N2 + 3H2 2NH3, the rate of reaction was... 

The balanced equation that we will use for this stoichiometry explanation is the recipe for the manufacture of ammonia (NH3). This reaction was so important that the chemist responsible for it, Fritz Haber, was awarded the Nobel Prize. Ammonia is a gateway step in the manufacture of fertilizers, and its manufacture was a giant step in solving the problem of providing food to a world population growing at an exponential rate. The equation for the Haber process is... 

One of the main goals of chemical kinetics is to understand the steps by which a reaction takes place. This series of steps is called the reaction mechanism. Understanding the mechanism allows us to find ways to facilitate the reaction. For example, the Haber process for the production of ammonia requires high temperatures to achieve commercially feasible reaction rates. However, even higher temperatures (and more cost) would be required without the use of iron oxide, which speeds up the reaction. 

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. 

Metal surfaces have long been appreciated for their catalytic properties the ability to alter the pathways of important chemical reactions so as to increase reaction rates and selectivity for desired products well above the rates occurring in a homogeneous phase. Catalytic metals are widely utilized in a variety of chemical processes, and have been critical for processes that have had worldwide energy and health benefits, such as the petroleum refining process for fuel production, the Haber-Bosch... 

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. 

Increasing pressure does not increase the number of particles, but it brings the particles closer together so collisions are more frequent. The Haber process, for example, uses high pressures to increase the rate of reaction of hydrogen and nitrogen to form ammonia. Lowering... 

Ammonia has been produced commercially from its component elements since 1909, when Fritz Haber first demonstrated the viability of this process. Bosch, Mittasch and co-workers discovered an excellent promoted Fe catalyst in 1909 that was composed of iron with aluminium oxide, calcium oxide and potassium oxide as promoters. Surprisingly, modem ammonia synthesis catalysts are nearly identical to that first promoted iron catalyst. The reaction is somewhat exothermic and is favoured at high pressures and low temperatures, although, to keep reaction rates high, moderate temperatures are generally used. Typical industrial reaction conditions for ammonia synthesis are 650-750 K and 150-300 atm. Given the technological importance of the... 

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... 

Operating at higher temperatures increases the reaction rate but reduces K (T) enormously because the reaction is exothermic and the variation of equilibrium constant with temperature is similar to that shown in Fig. 15.6. (For the Haber-Bosch process, = 60 moH dm at 227°C and at 0.02 mol dm at 527°C.)... 

Placed in its proper context, Haber s laboratory work, stimulated by academic polemics, demonstrated the possibilities for high pressure ammonia synthesis on an industrial scale. Unlike most other chemists whose activities straddled the worlds of academic and applied organic chemistry Haber opted for the uncertainties of physical chemistry. Not only did he show by clever manipulation of thermodynamics and kinetics that the reaction could be made to take place in yields that others considered impossible, but he also developed the process to the extent that the reaction rate was attractive to industrial chemists. Haber also knew that the only way to convince them would be to devise a continuous apparatus in which the valuable unreacted gases were recirculated to the converter. In July 1909 this apparatus, demonstrated at Karlsruhe, served its intended purpose. Haber s reward was a profit sharing arrangement with BASF, directorship of the new Kaiser Wilhelm-Institut fur Physikalische Chemie und Elektrochemie in Berlin (1912), and the Nobel Prize (1918). 

N2 and H2. In designing his process, Haber had to deal with a rapid decrease in the equilibrium constant with increasing temperature (Table 15.2). At temperatures sufficiently high to give a satisfactory reaction rate, the amount of ammonia formed was too small. The solution to this dilemma was to develop a catalyst that would produce a reasonably rapid approach to equilibrium at a suiBciently low temperature, so that the equilibrium constant remained reasonably large. The development of a suitable catalyst thus became the focus of Haber s research efforts. 

The Haber-Bosch catalytic process for production of ammonia is perhaps an invention that had the most dramatic impact on the human race (Ritter 2008). The inexpensive iron-based catalyst for ammonia synthesis, which replaced the original, more expensive osmium and uranium catalysts, made it possible to produce ammonia in a substantially effective manner. The objective here was not improvement in selectivity but higher reaction rates for rapid approach to the equilibrium conversion at the specified temperatme and pressme. Higher rates meant lower catalyst volume and smaller high-pressme reactors. The iron catalyst was improved by addition of several promoters such as alkali metals. In contrast to this simple single reaction case of ammonia synthesis, most organic reactions are complex with multiple pathways. 

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