Presentation on theme: "EQUILIBRIUM TIER 5 & TIER 6 TIER 5:Apply the concepts of kinetics and equilibrium to industrial processes TIER 6: Make connections between equilibrium,"— Presentation transcript:
EQUILIBRIUM TIER 5 & TIER 6 TIER 5:Apply the concepts of kinetics and equilibrium to industrial processes TIER 6: Make connections between equilibrium, thermodynamics, kinetics and gas laws.
HABER PROCESS HISTORY Early in the twentieth century, several chemists tried to make ammonia from atmospheric nitrogen. However this was nearly impossible until in 1909 Fitz Haber with help from Carl Bosch developed the Haber method. Before the Haber method, the synthetic ammonia needed for the production of nitric acid, a precursor to munitions was produced by the reaction of sodium nitrate with sulfuric acid. During World War I, the Allies had access to large amounts of sodium nitrate deposits in Chile (so called "Chile saltpetre") that belonged almost totally to British industries. As Germany lacked access to such readily available natural resources, the Haber process proved important to the German war effort. Fritz Haber’s contribution helped the German war effort, enabling the Germans to continue their war effort for another 4 years during World War I. However this contribution still did not matter with the rise of anti-semitism and because of his Jewish ancestry, he was expelled from the country in 1933.
The Haber method is also a major contributor to the production of fertilizer. It has transformed the world’s ability to produce food. So as controversial as the Haber method is to the German war effort, the contribution to the world’s food production can not be overlooked. Ammonia is also used in the production of many pharmaceuticals, in cleaners, as antimicrobial agents for food production and in the production of textiles. HABER PROCESS OTHER CONTRIBUTIONS
HABER PROCESS CHEMISTRY THE EQUATION IS: N 2 (g) + 3H 2 (g) 2NH 3 (g) ΔH = -93 kJ mol -1 THE REACTANTS AND PRODUCTS ARE ALL GASES THE FORWARD REACTION IS EXOTHERMIC THE RATIO OF NITROGEN TO HYDROGEN IS 1:3 THERE IS A TOTAL OF 4 MOLECULES ON THE REACTANT SIDE TO THE 2 MOLECULES FOR THE PRODUCT
HABER METHOD & LeCHATELIER’S PRINCIPLE THE REACTANTS AND PRODUCTS ARE ALL GASES: The reaction will proceed until it reaches equilibrium. The equilibrium constant is large enough(640 )that the production of ammonia should be favored but it takes to long for the production that it is not feasible THE FORWARD REACTION IS EXOTHERMIC: For the reaction to proceed to the right, it requires a low temperature THE RATIO OF NITROGEN TO HYDROGEN IS 1:3: The industrial process used to produce ammonia only produces about 15% yield with each pass through so the nitrogen and hydrogen is cycled back through repeatedly until it is used up. This is because of the temperature and pressure constraints as well the safety precautions used. THERE IS A TOTAL OF 4 MOLECULES ON THE REACTANT SIDE TO THE 2 MOLECULES FOR THE PRODUCT: This means that a high pressure favors the production of the ammonial
The catalyst The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency. The pressure The pressure varies from one manufacturing plant to another, but is always high. 200 atmospheres is the pressure usually used. Recycling At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%. The temperature In order to get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, °C isn't a low temperature! Rate considerations The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor. The compromise °C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time. Separating the ammonia When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled. The Haber process: the industrial procedure N 2 (g) + 3H 2 (g) 2NH 3 (g) ΔH = -93 kJ mol -1 STEPS: 1.Nitrogen and hydrogen are reacted 2.The ammonia is liquified 3.The unreacted gases are recycled
THE CONTACT METHOD PROCESS FOR SULFURIC ACID PRODUCTION THE CONTACT METHOD FOR THE PRODUCTION OF SULFURIC ACID TAKES PLACE IN A SERIES OF STEPS: 1.THE COMBUSTION OF SULFUR: S(s) + O 2 (g) SO 2 (g) 2. THE OXIDATION OF SULFUR DIOXIDE 2 SO 2 (g) + O 2 (g) 2SO 3 (g) ΔH = -196kJ mol THE COMBINATION OF SULFUR TRIOXIDE WITH WATER BY FIRST COMBINING THE SULFUR TRIOXIDE WITH CONCENTRATED SULFURIC ACID SO 3 (g) + H 2 SO 4 (l) H 2 S 2 O 7 (l) 2H 2 SO 4 (l)
THE CONTACT PROCESS & LeCHATELIER’S PRINCIPLE The temperature You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of sulphur trioxide in the equilibrium mixture. The forward reaction (the production of sulphur trioxide) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as low a temperature as possible. However, °C isn't a low temperature! Rate considerations The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much sulphur trioxide as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of sulphur trioxide if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor. The compromise °C is a compromise temperature producing a fairly high proportion of sulphur trioxide in the equilibrium mixture, but in a very short time. The pressure Notice that there are 3 molecules on the left-hand side of the equation, but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as high a pressure as possible. High pressures also increase the rate of the reaction. However, the reaction is done at pressures close to 2 atmospheres Economic considerations Even at these relatively low pressures, there is a 99.5% conversion of sulphur dioxide into sulphur trioxide. The very small improvement that you could achieve by increasing the pressure isn't worth the expense of producing those high pressures. The catalyst (V 2 O 5 ) Equilibrium considerations The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of sulphur trioxide in the equilibrium mixture. Its only function is to speed up the reaction. Rate considerations In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor.