Question

1) The equilibrium constant Kc of the reaction

I2 (g) ⇔ 2I (g)

is 5.6 × 10-12, at 500 K. In a system maintained at 500 K, the concentration of I2 is 0.020 mol / L and that of I is 2.0 × 10-8 mol / L. Is the reaction in balance? If not, in what sense does the reaction advance to reach equilibrium?


2) Given the concentrations below, how much is Q worth?

If Kc = 1.0, which side of the reaction is favored with this value of Q?

CO (g) + H2O (g) ⇄ CO2 (g) + H2 (g)

[CO (g)] = [H2O (g)] = 1.0 M

[CO2 (g)] = [H2 (g)] = 15 M

Answer 1

Step-by-step explanation:

1) Find the reaction quotient:

Qc = [I]² / [I₂]

Plug in the Kc value and the I₂ concentration.

Q = [2.0×10⁻⁸]² / (0.020)

Q = 2.0×10⁻¹⁴

Since Qc < Kc, the reaction moves to the right (favors the products).

2) Find the reaction quotient:

Qc = [CO₂] [H₂] / ([CO] [H₂O])

Qc = (15) (15) / ((1.0) (1.0))

Qc = 225

Since Qc > Kc, the reaction moves to the left (favors the reactants).