Question

Phosphorus pentachloride decomposes to phosphorus trichloride at high temperatures according to the reaction: PCl5 (g) ⇔ PCl3 (g) + Cl2 (g). At 250°C, 0.250 M PCl5 is added to a flask. If KC= 1.80, what are the equilibrium concentrations of each gas? A) [PCl5] = 1.25 M, [PCl3] = 0.474 M, [Cl2] = 0.474 M B) [PCl5] = 2.27 M, [PCl3] = 2.02 M, [Cl2] = 2.02 M C) [PCl5] = 0.0280 M, [PCl3] = 0.222 M, [Cl2] = 0.222 M D) [PCl5] = 1.80 M, [PCl3] = 1.80 M, [Cl2] = 1.80 M

Answer 1

C) [PCl5] = 0.0280 M, [PCl3] = 0.222 M, [Cl2] = 0.222 M

Step-by-step explanation:

Moles of [PCl₅] = 0.250 M

Considering the ICE table for the equilibrium as:

PCl₅ (g) ⇔ PCl₃ (g) + Cl₂ (g)

t = o 0.250

t = eq -x x x

--------------------------------------------- --------------------------

Moles at eq: 0.250-x x x

The expression for the equilibrium constant is:

`K_c=\frac {[PCl_3][Cl_2]}{[PCl_5]}=1.80`

So,

`(x^2)/(0.250-x)=1.80`

x = 0.222 M

Equilibrium concentrations :

[PCl₃] = [Cl₂] = 0.222 M

[PCl₅] = 0.250 - 0.222 = 0.280 M

Option C is correct.