The following equilibrium exists in a closed container: N2(g) + O2(g) = 2NO(g), ΔH = +181 kJ/mol. Which of the following perturbations would favor the formation of NO(g)? (practice MCAT question from The Princeton Review MCAT Science Workbook, 2009 edition)
increasing the pressure
decreasing the pressure
Increasing the temperature
decreasing the temperature
N2 + O2 + heat ==> 2NO
You go through the same reasoning as the last problem we did. You didn't have pressure in that one but the rule on pressure is an increase in P shifts the equilibrium to the side with fewer mols gas. Since you have 2 mols on the left and 2 mols on the right and increase or decrease in P will not affect the equilibrium. You want to increase NO so how can you do it.
increase the temperature ?
sure. Since it is an endothermic reaction, increasing T will shift it to the right and that will produce more NO
To determine which perturbation would favor the formation of NO(g) in this equilibrium, we can consider Le Chatelier's principle. Le Chatelier's principle states that if a system at equilibrium is subjected to a change, it will adjust to minimize the effect of that change and reach a new equilibrium position.
In this case, increasing the pressure would favor the side of the reaction with fewer moles of gas. By increasing the pressure, we are effectively increasing the concentration of the reactants (N2 and O2). Since there are fewer moles of reactants compared to the product (2 moles of NO), the reaction will shift to the right to form more NO to minimize the effect of the increased pressure. Therefore, increasing the pressure would favor the formation of NO(g).
On the other hand, decreasing the pressure would favor the side of the reaction with more moles of gas. By decreasing the pressure, we are effectively decreasing the concentration of the reactants (N2 and O2). Since there are fewer moles of reactants compared to the product, the reaction will shift to the left to form more reactants to minimize the effect of the decreased pressure. Therefore, decreasing the pressure would not favor the formation of NO(g).
Increasing the temperature would favor the endothermic reaction. Since the forward reaction is endothermic (ΔH = +181 kJ/mol), increasing the temperature would provide the necessary energy to favor the formation of NO. Therefore, increasing the temperature would favor the formation of NO(g).
Decreasing the temperature would favor the exothermic reaction. Since the forward reaction is endothermic, decreasing the temperature would decrease the energy available and thus favor the reverse reaction, which is the formation of N2 and O2. Therefore, decreasing the temperature would not favor the formation of NO(g).
In conclusion, increasing the pressure and increasing the temperature would favor the formation of NO(g) in this equilibrium.