For the reaction: H2(g) + I2(g) 2HI(g) , Kc = 12.3 at some temperature T. If [H2] = [I2] = [HI] = 3.21 x 10-3 M at this temperature, which one of the following statements is true?
1. ) Not enough information is given to answer the question.
2. ) The concentration of HI will rise as the system approaches equilibrium.
3. ) The concentrations of H2 and I2 will increase as the system approaches equilibrium.
4. ) The concentrations of H2 and HI will decrease as the system approaches equilibrium.
5. ) The system is at equilibrium, no change.
To answer this question, we need to compare the given concentrations of the reactants and products to the equilibrium constant (Kc).
The equilibrium constant expression for the given reaction is Kc = [HI]^2 / ([H2] * [I2]), where [HI], [H2], and [I2] represent the concentrations of HI, H2, and I2, respectively.
Given that [H2] = [I2] = [HI] = 3.21 x 10^-3 M, we can substitute these values into the equilibrium constant expression:
Kc = (3.21 x 10^-3)^2 / ((3.21 x 10^-3) * (3.21 x 10^-3))
Simplifying, we get:
Kc = 1.037
Since the calculated equilibrium constant (Kc = 1.037) does not equal the given equilibrium constant (Kc = 12.3), we can conclude that the system is not at equilibrium.
Therefore, the correct statement is option 3: The concentrations of H2 and I2 will increase as the system approaches equilibrium.
To determine which statement is true, let's compare the given concentrations with the equilibrium constant (Kc = 12.3). The equilibrium constant expression for the reaction is:
Kc = [HI]² / ([H2] * [I2])
Plugging in the values:
12.3 = ([3.21 x 10-3 M])² / ([3.21 x 10-3 M] * [3.21 x 10-3 M])
Simplifying:
12.3 = [3.21 x 10-3 M] / [3.21 x 10-3 M]
The concentration of HI cancels out, leaving only a constant value. This means that the concentration of HI is already at equilibrium and will not change as the system approaches equilibrium. Therefore, statement 2, 3, 4, and 5 are incorrect.
The correct answer is: 1. Not enough information is given to answer the question.