Consider an exothermic chemical reaction
A (g) + B (g) ⇌ C (g) + 2D (s)
1.00 mole of A and 1.00 mole of B are placed in a 2.00 liter container. After equilibrium has
been established, 0.10 mole of C is present in the container. (1pt)
a) Calculate the equilibrium constant for the reaction.
b) Based on your answer to a), does the equilibrium favor formation of reactants or
products?
a) The equilibrium constant for the reaction is K = [C]^2/[A][B] = 0.01.
b) Based on the answer to a), the equilibrium favors formation of reactants.
To calculate the equilibrium constant for the given exothermic chemical reaction, we need to use the equilibrium expression and the concentrations of the reactants and products at equilibrium.
The equilibrium constant expression for the reaction is given by:
Kc = [C]^c [D]^d / [A]^a [B]^b
Where [C], [D], [A], and [B] represent the molar concentrations of C, D, A, and B respectively, and a, b, c, and d are the stoichiometric coefficients of the reactants and products in the balanced equation.
In the given reaction:
A (g) + B (g) ⇌ C (g) + 2D (s)
a = 1 (stoichiometric coefficient of A)
b = 1 (stoichiometric coefficient of B)
c = 1 (stoichiometric coefficient of C)
d = 2 (stoichiometric coefficient of D)
The given equilibrium concentrations are:
[C] = 0.10 mole (given)
[A] = 1.00 mole (initially)
[B] = 1.00 mole (initially)
Substituting the values into the equilibrium constant expression, we have:
Kc = (0.10)^1 (1.00)^1 / (1.00)^1 (1.00)^1
Kc = 0.10
Therefore, the equilibrium constant for the reaction is 0.10.
Now moving on to part (b) of the question.
Since the equilibrium constant (Kc) is 0.10, we can say that the equilibrium does not favor the formation of products. Typically, if the equilibrium constant is less than 1, it indicates that the concentration of products is lower than the concentration of reactants at equilibrium. In this case, the equilibrium favors the formation of reactants rather than products.
To calculate the equilibrium constant for the reaction, you can use the equation:
Kc = ([C]^c * [D]^d) / ([A]^a * [B]^b)
Where [C], [D], [A], and [B] represents the concentrations of each species at equilibrium, and a, b, c, and d represent the stoichiometric coefficients of each species in the balanced equation.
In this case, at equilibrium, the concentration of C is 0.10 mol/L (since it is given in the problem). The concentration of D can be determined by multiplying the molar ratio of C to D (2:1) with the concentration of C:
[D] = 2 * [C] = 2 * 0.10 mol/L = 0.20 mol/L
The concentrations of A and B at equilibrium can be calculated by subtracting the moles of C and D from their initial moles, respectively:
[A] = 1.00 mol - 0.10 mol = 0.90 mol
[B] = 1.00 mol - 0.20 mol = 0.80 mol
Now, you can substitute these values into the equilibrium constant expression:
Kc = ([C]^c * [D]^d) / ([A]^a * [B]^b)
= (0.10^1 * 0.20^2) / (0.90^1 * 0.80^1)
= 0.04 / 0.72
= 0.0556
So, the equilibrium constant for the reaction is approximately 0.0556.
Since the equilibrium constant (Kc) is less than 1, it indicates that the equilibrium favors the formation of reactants rather than products.