Favourable Reaction Conditions

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Favourable Reaction Conditions One can use Le Châtelier's principle to explain the most favourable reaction conditions for the production of nitrogen (II) oxide. By decreasing the pressure, temperature or by adding/removing different components of the reaction, one could manipulate the equilibrium creating the most favourable reaction conditions. By decreasing the pressure the equilibrium will shift by producing more molecules in order to increase the pressure of the reaction again. The equilibrium will move in such a way that the pressure increases again. Due to the fact that the ratio of the molecules is 9:10 in the reaction, by decreasing the pressure the position of equilibrium will move towards the right-hand side of the reaction. Lowering the pressure will favour the higher volume on the right side (ratio; 9:10) and therefore as a result will shift to the right. To do this the system would have to create more products and because nitrogen (II) oxide is a product, more nitrogen (II) oxide would be created. Therefore, this would result in the increased production of nitrogen (II) oxide. Along with decreasing pressure, decreasing temperature can also maximize the production of nitrogen (II) oxide. In dynamic equilibrium decreasing the temperature of a system favors the exothermic reaction and as a result the system counteracts the change that has been inflicted by producing more heat. Decreasing the temperature would shift the equilibrium to the right due to the fact that the forward reaction is exothermic. In turn, this would favour the production of nitrogen (II) oxide. According to Le Châtelier's principle, adding something will cause the system to shift away from it. That being said, by adding O2 (g) or NH3 (g) the dynamic equilibrium would disturbed because of the changing conditions and the position of equilibrium would move to counteract the

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