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.
