# Question #0ef5a

Jan 26, 2015

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 ${K}_{e q}$, 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 ${K}_{e q}$, 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 ${K}_{e q}$, the reaction is at equilibrium and no shift will take place.

${Q}_{c} = \frac{\left[B\right]}{\left[A\right]} = \left(\text{0.100 mol/L")/("0.020 mol/L}\right) = 5.0$

Notice that ${Q}_{c}$ is smaller than ${K}_{e q}$, which is said to be 10.0. This means that the reaction will shift to the right and favor the formation of more product.