For the equilibrium reaction
Co(g)+H2O(g)<=>CO2(g)+H2(g)
the Keq value at 690°C is 10.0. A mixture of 0.300 mol of CO, 0.300 mol of H2O, 0.500 mol of CO2 and 0.500 mol of H2 is placed in a 1.0 L flask.
a) Write the Keq expression for this reaction and use the concentrations above to calculate a trial value of the Keq. Does this value match the given Keq ? Is the reaction at equilibrium ? Explain.
b)Calculate the equilibrium concentrations of all four species.
I worked this last night for you.
a) The equilibrium constant expression (Keq) for a reaction in terms of concentrations is determined by writing the concentrations of the products divided by the concentrations of the reactants, each raised to the power of their respective stoichiometric coefficients. For the reaction given:
Keq = [CO2]^1 * [H2]^1 / [CO]^1 * [H2O]^1
Substituting the given concentrations into the expression:
Keq (trial) = (0.500)^1 * (0.500)^1 / (0.300)^1 * (0.300)^1
Keq (trial) = 0.25 / 0.09
Keq (trial) ≈ 2.78
The calculated trial value of Keq is approximately 2.78, which does not match the given Keq value of 10.0. Therefore, the reaction is not at equilibrium.
b) To calculate the equilibrium concentrations, we need to use an approach called the ICE (Initial-Change-Equilibrium) method. We start with the initial concentrations given:
[CO] = 0.300 M
[H2O] = 0.300 M
[CO2] = 0.500 M
[H2] = 0.500 M
Since the reaction is not at equilibrium, we need to calculate the changes in concentration. Let's represent the change in concentration as x, assuming it is the same for CO, H2O, CO2, and H2.
[CO] changes to [CO] + x
[H2O] changes to [H2O] + x
[CO2] changes to [CO2] - x
[H2] changes to [H2] - x
Now let's substitute these values into the Keq expression:
Keq = ([CO2] - x) * ([H2] - x) / ([CO] + x) * ([H2O] + x)
To simplify calculations, we neglect the x term because it is small compared to the initial concentrations. Therefore, the expression becomes:
Keq ≈ [CO2] * [H2] / [CO] * [H2O]
Now we can substitute the given concentrations into the Keq expression:
Keq = (0.500) * (0.500) / (0.300) * (0.300)
Keq ≈ 0.833
The calculated value of Keq is approximately 0.833, which does not match the given Keq value of 10.0. Hence, the reaction is not at equilibrium.