Step-by-step explanation:
Given the following reaction:
N₂ (g) + 3 H₂ (g) <----> 2 NH₃ (g) K = 9.60
We know the initial concentrations of the species.
[N₂] = [H₂] = 0.100 M [NH₃] = 0.352 M
The expression for the equilibrium constant is:
k = [NH₃]²/([N₂]*[H₂]³)
We can replace the given values of the concentrations to get the quotient of the reaction and then compare it to the value of the equilibrium constant to determine whether if the reaction is at equilbrium or not.
Q = [NH₃]²/([N₂]*[H₂]³)
Q = (0.352)²/((0.100*(0.100)³)
Q = 1239
Since the Quotient of the reaction is not equal to the equilibrium constant we can say that the reaction is not at equilibrium.
Q > k ---> not at equilibrium
To reach the equilibrium we must reduce the value of Q. Q (or k) is the concentration of the products over the concentration of the reactants (all of them powered to their coefficients). So if we want to reduce that number we have two options.
Q = [products]^n/[reactants]m
To reduce the value of Q the numerator must be smaller or we can increase the denominator. That means that we can reduce the products or increase the reactants. We have to shift the equilibrium to the left.
Ammonia must be consumed and nitrogen and hydrogen gas must be produced.
Answer: c) More nitrogen and hydrogen must form to achieve equilibrium.