7.5 Quantitative Changes in Equilibrium Systems Day 2 (more complex questions)

Let's consider a simple example to calculate the reaction quotient, Q, for the following hypothetical reaction:
N2(g) + 3H2(g) ⇌ 2NH3(g)
Suppose you have the following concentrations at a certain point in the reaction:
[N2] = 0.2 M
[H2] = 0.5 M
[NH3] = 0.3 M
Now, let's calculate Q for this reaction using the given concentrations:
Q = [NH3]^2 / [N2][H2]
Q = (0.3 M)^2 / (0.2 M)(0.5 M)
Q = 0.09 M^2 / 0.1 M^2
Q = 0.9
So, at this particular moment in the reaction, Q is 0.9. Now, we can compare this value to the equilibrium constant, K, to determine the direction in which the reaction will proceed:
If Q is less than K, the reaction will shift to the right (toward the products) to reach equilibrium.
If Q is greater than K, the reaction will shift to the left (toward the reactants) to reach equilibrium.
If Q = K, the reaction is at equilibrium.
Suppose the equilibrium constant K for this reaction is 5.0. Since Q (0.9) is less than K (5.0), the reaction will shift to the right, meaning that more product, ammonia (NH3), will be formed until Q reaches the equilibrium constant K.