H2 (g) + I2 (g) 2 HI (g)
If the initial concentrations of H2 and I2 are 1.0 M and the initial concentration of HI is 0.5 M (Kc = 54.3 at 430oC).
(a) Is the reaction at equilibrium?
(b) If not, which way will the reaction proceed?
Let K=[HI]²/([H2][I2])
The given reaction is in equilibrium at 430°C, if K=Kc=54.3
Substitute given concentrations to get
K=[HI]²/([H2][I2])
=0.5²/(1*1)
=0.25 < 54.3
Therefore the reaction is not in equilibrium and will proceed in the forward direction until K=Kc.
To determine whether the reaction is at equilibrium, we need to compare the calculated reaction quotient, Qc, with the equilibrium constant, Kc.
(a) To calculate Qc, we use the formula Qc = [HI]^2 / ([H2] * [I2]). Given that the initial concentration of HI is 0.5 M, and the initial concentrations of H2 and I2 are 1.0 M, we can substitute these values into the formula:
Qc = (0.5^2) / (1.0 * 1.0) = 0.25 / 1 = 0.25
Next, we compare Qc with Kc. If Qc = Kc, the reaction is at equilibrium. If Qc < Kc, the reaction will proceed in the forward direction to reach equilibrium. And if Qc > Kc, the reaction will proceed in the reverse direction to reach equilibrium.
Given that Kc = 54.3, and Qc = 0.25, we can conclude that Qc < Kc. Therefore, the reaction is not at equilibrium.
(b) Since Qc < Kc, the reaction will proceed in the forward direction to reach equilibrium. This means that more of the products, HI, will be formed from the reactants, H2 and I2.
So overall, the reaction is not at equilibrium (a), and it will proceed in the forward direction (b).