How will the equilibrium shift for this reaction if you change the temperature and pressure? Describe this by using your knowledge of Le Chatelier's principle and the Haber Bosch method for this reaction:  N2(g) + 3H2 (g) --> 2NH3(g)

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The Haber-Bosch process for producing ammonia from atmospheric nitrogen and hydrogen was invented in Germany during World War I, initially to produce ammonia for military explosives. It's now an important method of producing nitrogen fertilizer for agriculture.

According to Le Chatelier's Principle, if a system at equilibrium is disturbed the...

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The Haber-Bosch process for producing ammonia from atmospheric nitrogen and hydrogen was invented in Germany during World War I, initially to produce ammonia for military explosives. It's now an important method of producing nitrogen fertilizer for agriculture.

According to Le Chatelier's Principle, if a system at equilibrium is disturbed the position of the equilibrium will shift so as to counteract the affect of the change. For example, if the temperature is increased the equilibrium of an exothermic reaction will shift to the left, the direction that favors reactants, because the reverse reaction absorbs heat. The equilibrium of an endothermic reaction will shift to the right, or produce more product, because the forward direction absorbs heat.

The reaction`3 H_2 + N_2 -> 2 NH_3` is exothermic. Therefore a decrease in temperature will result in a shift in equilibrium that favors products. However, this reaction proceeds slowly at room temperature so lowering temperature to shift equilibrium was not a good option. The process is carried out at about 400 *C as that maximizes the trade-off between speed and yield.

The effect of a pressure change is related to the moles of gaseous reactants and products. An increase in pressure shifts the equilibrium to the direction that produces fewer gas particles and a decrease in pressure shifts the equilibrium to the direction that produces more gas molecules.

The reactants are gases and ammonia is produced as a gas at the reaction temperature so there's a ratio of 4 moles gaseous reactant to 2 moles gaseous product. This means that raising the pressure will shift the equilibrium in the direction of the product and increase the yield. High pressure proved to be a challenge in the development of the Haber Process as vessels and pipes had to be very strong to withstand it. The pressure that's used, about 200 atm, is the optimum safe pressure for the best yield.

Raising the temperature (which didn't shift the equilibrium in the desired direction) and raising the pressure didn't produce an efficient enough process. Another application of LeChatelier's principle, condensing and removing the ammonia as it formed, shifted the equilibrium to produce more product and contributed to the successful production of ammonia on an industrial scale. The other important factor was the use of a catalyst, which didn't shift the equilibrium but did speed up the reaction.

If you're interested in learning more about the history of ammonia use and production, I recommend the book The Alchemy of Air by Thomas Hager, published in 2008.

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