Question

For the following reaction, Kc = 15 at 700 K. 2 NO(g) + Cl2(g) ⇄ 2 NOCl(g) If we have [NO] = 0.15 M, [Cl2] = 0.15 M, [NOCl] = 0.40 M at 700 K, what will happen? Group of answer choices The equilibrium will not shift. The equilibrium will shift to the left, but will use up only part of the NOCl. The equilibrium will shift to the right, but will use up only part of the NO and Cl2. The equilibrium will shift to the right until all the reactants are used up. The equilibrium will shift to the left until all the NOCl is used up.

Answer 1

The equilibrium will shift to the right, but will use up only part of the NO and Cl2.

Step-by-step explanation:

The given reaction is 2 NO(g) + Cl2(g) ⇄ 2 NOCl(g), and the equilibrium constant (Kc) is 15 at 700 K. The initial concentrations are [NO] = 0.15 M, [Cl2] = 0.15 M, and [NOCl] = 0.40 M. To determine the direction of the equilibrium shift, we compare the initial concentrations to the equilibrium concentrations.

In this case, the concentrations of NO and Cl2 are greater than the equilibrium concentrations, while the concentration of NOCl is less than the equilibrium concentration. Based on Le Chatelier's principle, the equilibrium will shift to the right to counteract the decrease in NOCl concentration and increase the concentrations of NO and Cl2.

Therefore, the correct answer is: The equilibrium will shift to the right, but will use up only part of the NO and Cl2.