The equilibrium will shift to the right, favoring the formation of more product.
So, in order to determine whether or not a reaction is at equilibrium, you must calculate the reaction quotient, or #Q_c#.
#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.
If #Q_c# is equal to #K_(eq)#, the reaction is at equilibrium and no shift will take place.
Notice that #Q_c# is smaller than #K_(eq)#, which is said to be 10.0. This means that the reaction will shift to the right and favor the formation of more product.
