The reaction A (aq) ⇌ B (aq) + C (aq) has an equilibrium constant given by:
Kc = [B]eq [C]eq / [A]eq
Where [X]eq is the equilibrium concentration of X in mol dm⁻³.
If a 2.0 mol dm⁻³ solution of A is allowed to reach equilibrium, at which 1.2 mol dm⁻³ of A remains, find the equilibrium concentrations of B and C in mol dm⁻³. Give your answer to 1 significant figure.
Answer (1)
To solve this problem, we need to find the equilibrium concentrations of B and C given the initial concentration of A and the amount of A remaining at equilibrium.
Identify the initial and equilibrium concentrations:
Initially, we have a 2.0 mol dm⁻³ solution of A, and no B or C are present.
At equilibrium, 1.2 mol dm⁻³ of A remains, so the change in concentration of A is 2.0 - 1.2 = 0.8 mol dm⁻³.
Determine the change in concentrations:
For the reaction A \(aq \leftrightarrow B a q + C a q ), as A decreases by 0.8 mol dm⁻³, both B and C increase by the same amount, since they are produced in a 1:1 ratio.
Therefore, [ B ] e q = 0.8 mol dm⁻³ and [ C ] e q = 0.8 mol dm⁻³.
Calculate the equilibrium concentrations:
At equilibrium: [ A ] e q = 1.2 mol dm − 3 , [ B ] e q = 0.8 mol dm − 3 , [ C ] e q = 0.8 mol dm − 3 .
Thus, the equilibrium concentrations of B and C are both 0.8 mol dm⁻³, rounded to 1 significant figure, the answer is 0.8 mol dm⁻³ for each. This is the final state of the system after it has reached equilibrium under the conditions provided.
