Question

1) Consider the equilibrium reaction

N₂(g) + 3 H₂(g) ↔ 2 NH₃(g) + 91.8 kJ

Using Le Chatelier’s Principle, Name two specific ways this equilibrium can be shifted to the right to increase production of NH3.

2) In the reaction below,

NH₃(g) + H₂O(l) ↔ NH₂ - + H₃O +

The equilibrium constant is 1 x 10 ^-34. Is this reaction likely to take place? Explain your answer.

Answer 1

1) We see that number of moles of gas in the reactant side (3+1 = 4 moles) is more than the number of moles of gas in the product side (2 moles). So if the reaction moves in the forward direction, the pressure of the system is being reduced.

Hence, if we INCREASE the PRESSURE of the reaction, the reaction will move in the forward direction to reduce the pressure thus producing more NH3

It is given that delta H (Enthalphy change) for the reaction is +91.8 kJ. The positive sign signifies that the reaction is endothermic. Hence energy is being consumed by the reacion in the forward direction.

Therefore, if we DECREASE the TEMPERATURE of the reaction the equilibrium could be shifted to the right increasing theproduction of NH3

2)

NH₃(g) + H₂O(l) ↔ NH₂ - + H₃O +

Equilibrium Constant (K) = 10^-34

K = [NH2-] [H3O+] / [NH3] [H2O]

Since K << 1,

[NH2-] and [H3O+] << [NH3] and [H2O]

The reaction will take place but it will be negligible since the value of K is very small.