Answer:
Q<K so N2(g) +3H2(g) → 2NH3(g)
Explanation:
Step 1: Data given
For the balanced equation N2(g)+3H2(g)⇌2NH3(g) We have the following data:
⇒ Reaction quotient Q = 3.56 *10^-4
Where Q is defined as Q = [NH3]^2/{[N2]*[H2]^3]
⇒ Equilibrium constant K = 6.02 * 10^-2
If we compare Q and K, there are 3 options
⇒ Q > K : The reaction favors the reactants. This means that in the Q equation, the ratio of the concentration or pressure of the products) to the concentration or pressure of the reactants is larger than that for K.
This means that more products are present than there would be at equilibrium.
Following Le chatelier's principle reactions always tend toward equilibrium, this means the reaction will produce more reactants from the excess products, and will therefore cause the system to shift to the LEFT.
Doing this, it will allow the system to reach equilibrium.
⇒ Q = K : The reaction is already at equilibrium. There is no tendency to form more reactants or more products at this point. No side is favored and no shift will occur.
⇒ Q < K : The reaction favors the products. The ratio of products to reactants is <u>less</u> than that for the system at equilibrium—the concentration or the pressure of the reactants is greater than the concentration or pressure of the products.
Following Le chatelier's principle reactions always tend toward equilibrium, the system shifts to the RIGHT to make more products.
In this case K > Q since 6.02*10^-2 > 3.56*10^-4
This means that forward reaction will be favored or the reaction will shift to the right
N2(g) +3H2(g) → 2NH3(g)
As a result, the reaction will consume nitrogen gas (N2) and hydrogen gas (H2), and produce more ammonia (NH3).
The reaction will proceed in this direction until equilibrium is established
