During the World War I, Germany was dependent on sodium nitrate deposits in Chile for the nitrogen-containing compounds needed to make explosives. Seeking to free Germany from dependence on Chile, a German chemist named Fritz Haber examined the direct combination reaction between nitrogen and hydrogen, a reaction that forms ammonia.If we decrease the concentration of ammonia in the system, this decreased concentration of ammonia is a stress externally applied to the above system. According to LeChatelier's Principle the system will favor the process that replaces the ammonia that we siphoned off. The forward reaction or process produces ammonia. On other hand, if we increase the concentration of Nitrogen that would be a stress. The system will have to respond in order to undo the stress of increased Hydrogen. The forward process uses up Hydrogen and Nitrogen while the reverse process increases the presence of Nitrogen and Hydrogen.The industrial conditions for producing ammonia the temperature must be 450ºC to 500ºC. The forward reaction is exothermic. If we remove heat as a product will result in the equilibrium mixture making richer ammonia. Since we want ammonia from the Haber process, the reaction conducted at 450ºC. Because all reactions go faster if the temperature is raised. Reversible reactions, such as the Haber process, rasing the temperature will make the equilibrium mixture richer in nitrogen and hydrogen because forming these from ammonia takes heat in. If we cool the reaction down the proportion of ammonia will increase but the rate of production will decrease. With a reversible reaction, a catalyst which increases the rate will increase the rate of both the forward and the backward reaction. This is useful because the catalyst will, cause the reaction mixture to reach its equilibrium composition more quickly. The catalyst will not change the equilibrium composition of the substance.On its own, the reaction reached equilibrium long before any sizable amount of ammonia was produced. Haber and his team tried to maximize the equilibrium yield of ammonia. According to Le Chatelier's Principle, by increasing the temperature and pressure, the reaction tends to shift to the fewer molecule side, therefore the amount of the equilibrium yield of ammonia was maximized.Harber was able to substantially increase the ammonia yield by analyzing the its equilibrium mixture. The equilibrium expression for these reactions isN2(g) + H2(g) - NH4(g)+ HeatThe equilibrium mixture of N2 , H2 , and NH4 depends on ht initial concentration of each component. Harber and his team determined the equilibrium constant for this reaction by starting with various mixtures of the three components and allowing them to reach equilibrium at a specific temperature and pressure. They them sifted the equilibrium to produce the maximum amount of ammonia.Bibliography:Internet resources (http://members.aol.com/profchm/le_chat.html) (http://haberchemistry.tripod.com/) Prentice Hall Chemistry
