To determine the equilibrium concentration of CIF3 (g) after the reaction has occurred, we must first consider the stoichiometry of the reaction. From the balanced equation, we see that 1 mol of Cl reacts with 3 mol of F2 to produce 2 mol of CIF3.
Let's assume the initial concentration of Cl is [Cl]0, and the initial concentration of F2 is [F2]0. Since the reaction uses up 1 mol of Cl for every 3 mol of F2, the change in concentration of Cl will be -3x (where x is the change in concentration of F2).
Since the ratio Cl:F2 is 1:3, the change in concentration of F2 will be -x.
The equilibrium concentration of Cl will be [Cl]0 - 3x, and the equilibrium concentration of F2 will be [F2]0 - x. The equilibrium concentration of CIF3 will be 2x.
From the given equilibrium constant, we know:
Kc = [CIF3]^2 / ([Cl]*[F2]^3)
Substitute the equilibrium concentrations:
-0.024 = (2x)^2 / ([Cl]0 - 3x)*([F2]0 - x)^3
Since it seems difficult to solve this algebraic equation directly, you may need to assume an initial concentration of Cl and F2 to calculate A[CIF 3]/ At.
Consider the reaction Cl (g) + 3F2(g) → 2CIF3 (g),
4/Cl]/At = -0.024 М/в.
Determine A[CIF 3]/ At
1 answer
1 answer