Answer:
[HI] = 0.784M
[I₂] = 0.108M
[H₂] = 0.108M
Explanation:
Based on the equilibrium reaction:
H₂(g) + I₂(g) ⇄ 2HI(g)
The equilibrium constant, K, is:
K = 52.8 = [HI]² / [H₂] [I₂]
Where [] are equilibrium concentrations of each gas.
As initial concentration of HI is 1.00M, the equilibrium concentrations of the gases is:
[HI] = 1.00M - 2X
[I₂] = X
[H₂] = X
Replacing:
52.8 = [1.00-2X]² / [X] [X]
52.8X² = 4X² - 4X + 1
0 = -48.8X² - 4X + 1
Solving for X:
X = -0.1899M. False solution, there is no negative concentrations
X = 0.108M. Right solution.
Replacing, equilibrium concentrations are:
[HI] = 1.00M - 2*0.108M
