I keep getting the wrong answer for this one can someone help. The initial concentration for the compounds involved in the reaction displayed were determined to be [NOCl(g)] = 0.1738 mol/L, [NO(g)] = 0.1991 mol/L, [Cl2(g)] = 0.000005938 mol/L. Calculate the value of the equilibrium constant (Kc) at 25.00 °C if the equilibrium concentration of NO(g) was 0.1991 mol/L.
2NOCl(g) = 2NO(g)+Cl2(g)
I can't possibly find your error without your work. It is an odd problem that has initial and final concentrations the same. Are you certain this is the problem?
keq= (.1991)^2 *(5.938E-7)/(.1738)^2
To calculate the value of the equilibrium constant (Kc) for the given reaction, you need to determine the equilibrium concentrations of all the compounds involved and then use the equation relating the equilibrium constant to the concentrations.
In this case, you are given the initial concentrations of [NOCl(g)], [NO(g)], and [Cl2(g)], and you need to calculate the equilibrium concentration of [NO(g)].
Since the stoichiometric ratio between NOCl and NO is 2:2, the equilibrium concentration of NO(g) will also be 0.1991 mol/L, as given.
Now, you can use these equilibrium concentrations in the equation for the equilibrium constant (Kc):
Kc = ([NO(g)]^2 * [Cl2(g)]) / [NOCl(g)]^2
Plugging in the values:
Kc = (0.1991^2 * 0.000005938) / (0.1738^2)
Kc = 0.000002965 / 0.03012044
Kc ≈ 9.83 x 10^-5
Therefore, the value of the equilibrium constant (Kc) at 25.00 °C is approximately 9.83 x 10^-5.