Haber synthesis

If the percent yield of a reaction is already known, we can calculate how much of a product to expect from a synthesis that uses a known amount of starting material. For example, the Haber synthesis of ammonia stops when 13% of the starting materials have formed products. Knowing this, how much ammonia could an industrial producer expect to make from 2.0 metric tons of molecular hydrogen First, calculate the theoretical yield ... 

C04-0122. The following diagram represents a small portion of a reaction vessel that contains the starting materials for the Haber synthesis of ammonia ... 

C04-0140. In the Haber synthesis of ammonia, N2 and H2 react at high temperature, but they never react... 

C05-0139. The Haber synthesis of ammonia occurs in the gas phase at high temperature (400 to 500 °C) and... 

C15-0055. A chemist is studying the rate of the Haber synthesis N2 + 3 H2 2 NH3 Starting with a... 

C15-0122. Explain in molecular terms why the Haber synthesis cannot proceed in a single-step elementary reaction. 

The Haber synthesis for the conversion of N2 into NH3 is outlined below. The process uses N2 from the... 

Subsequent studies have failed to support the carbide theory, and it is now generally accepted that carbides of the type proposed by Craxford play little or no part in the Fischer-Tropsch synthesis (86, 87). It has, however, recently been suggested, by analogy with the mechanism proposed for the Haber synthesis of ammonia, that carbides formed by dissociative absorption of carbon monoxide would be expected to be readily hydrogenated and could therefore be of importance in Fischer-Tropsch synthesis over heterogeneous catalyst (88). 

Haber synthesis of ammonia from N2 and H2 Hydrogenation of vegetable oils Oxidation of hydrocarbons and CO and reduc-... 

To illustrate how an equilibrium mixture is affected by a change in pressure as a result of a change in the volume, let s return to the Haber synthesis of ammonia. The balanced equation for the reaction has 4 mol of gas on the left side of the equation and 2 mol on the right side ... 

By contrast, a change in temperature nearly always changes the value of the equilibrium constant. For the Haber synthesis of ammonia, which is an exothermic reaction, the equilibrium constant Kc decreases by a factor of 1011 over the temperature range 300-1000 K (Figure 13.12). 

Because N204 is colorless and N02 has a brown color, the effect of temperature on the N204-N02 equilibrium is readily apparent from the color of the mixture (Figure 13.13). For an exothermic reaction, such as the Haber synthesis of NH3, heat is absorbed by net reaction in the reverse direction, so Kc decreases with increasing temperature. 

Because the free-energy change for any process at constant temperature and pressure is AG = AH — TAS, we can calculate the standard free-energy change AG° for a reaction from the standard enthalpy change AH0 and the standard entropy change AS°. Consider again the Haber synthesis of ammonia ... 

Take the Haber synthesis, for example. The hypothetical process is... 

Table 5.1 shows an application of XPS to the study of the promoted iron catalyst used in the Haber synthesis of ammonia. The sizes of the various electron intensity peaks allows a modest level of quantitative analysis. This catalyst is prepared by sintering an iron oxide, such as magnetite (Fe304) with small amounts of potassium nitrate, calcium carbonate, aluminium oxide and other trace elements at about 1900 K. The unreduced solid produced on cooling is a mixture of oxides. On exposure to the nitrogen-hydrogen reactant gas mixture in the Haber process, the catalyst is converted to its operative, reduced form containing metallic iron. As shown in Table 5.1, the elemental components of the catalyst exhibit surface enrichment or depletion, and the extent of this differs between unreduced and reduced forms. 

However, if a reaction is exothermic (AH < 0), Le Chatelier s principle states that increasing the temperature during the reaction will push the thermodynamic equilibrium towards reactants. The classic example of this dilemma is seen in the Haber synthesis of ammonia, equation (1.17). 

Fritz Haber Synthesis of ammonia from its elements... 

F. Haber s catalytic synthesis of NH3 developed in collaboration with C. Bosch into a large-scale industrial process by 1913. (Haber was awarded (he 1918 Nobel Prize in Chemistry for (he synthesis of ammonia from its elements Bosch shared the 1931 Nobel Prize for contributions to the invention and development of chemical high-pressure methods , the Haber synthesis of NH3 being (he first high-pressure industrial process.)... 

This reaction is quite slow under normal conditions, but the rate can be greatly increased by the presence of a catalyst such as nickel or platinum. As in the Haber synthesis of ammonia (see Section 13.6), the main function of the catalyst is to weaken the H—H bond and facilitate the reaction. 

As a protection gas for stainless steels during heat treatment (1050°C), a mixture of 75% hydrogen and 25% nitrogen is very much used. It is prepared by cracking of ammonia at 900°C and atmospheric pressure with a nickel catalyst present. The process thus compels the Haber synthesis to go backwards. 

Fritz Haber, a German-born chemist, solved the problem. The Haber synthesis, as it is now called, is based on the equilibrium... 

The Haber synthesis of ammonia is the major industrial method for fixing nitrogen, but it requires high temperatures and pressures and is, therefore, expensive. 

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