Consider the following reaction, initially at equilibrium:?
N2(g) + 3H2(g) <—> 2NH3(g) + 92 kJ/mol
1. Which way does the equilibrium shift if you increase pressure on the whole system?
2. What happens to [H2] if you decrease [NH3]?
3. Which way does the equilibrium shift if you increase the temperature on the whole system?
- bee killerLv 71 month agoFavorite Answer
1. the reaction will shift to the products. Pressure will have a greater effect on the side of the reaction with the greatest number of moles of gasses,
2.H2 will decrease. By decreasing the NH3 the reverse reaction stops/decreases while the forward reaction continues and the H2 gets used up.
3. Increasing temp always favors endothermic reactions. In this case it is the reverse reaction.
- pisgahchemistLv 71 month ago
Le Chatelier's principle....
N2(g) + 3H2(g) <==> 2NH3(g) ................. ΔH = +92 kJ
1. Increase P ... Since P is always a dependent variable, you can increase P by increasing the temperature, or by decreasing the volume, or by increasing the number of molecules. Changing the temperature or the number of molecules have other ramifications, so look only at the volume change. When the pressure is increased, the system will shift in a direction which will decrease the pressure according to Le Chatelier's principle, and that will be toward the side with fewer molecules. Therefore, decreasing the volume and increasing the pressure shifts the system to the right.
2. Decrease [NH3] ... Removing ammonia shifts the system to the right to make more ammonia. This is one of the things that Haber and Bosch did to increase the yield of ammonia. Therefore, removing ammonia will decrease the concentration of hydrogen.
3. Increase T ... In an endothermic reaction we can pretend that heat is a reactant, so when more heat is added, the system shifts to the right to "use up" the added heat. And the value of Kc will increase.