Your reaction is at equilibrium.
In order to determine whether or not a reaction is at equilibrium, you must calculate the reaction quotient, or "Q"_c. The value of "Q"_c will tell you in which direction the reaction will progress if equilibrium has not yet been reached.
Q_c expressed the ratio of products to reactants at a given instant. If the value you obtain for Q_c is smaller than K_(eq), the equilibrium constant, there are more reactants than products, which will cause the equilibrium to shift to the right, favoring the formation of more products.
If Q_c is bigger than K_(eq), there are more products than reactants, which will cause the equilibrium to shift ot the left, favoring the formation of more reactants.
Finally, if Q_c is equal to K_(eq), the reaction is at equilibrium and no shift will take place.
So, let's calculate Q_c and compare it with K_(eq). Remember that you must express the concentrations of the species in "mol/L", not "mmol/L"
Q_c = ([K] * [L])/([J]) = (1.10 * 10^(-3) M * 125 * 10^(-3)M)/(11.0 * 10^(-3)M)
Q_c = (0.125 * 0.0011)/(0.011) = 0.0125
As you can see, Q_c is equal to K_(eq), which means the reaction is at equilibrium and no shift will take place.
