Question

Consider the equilibrium $$ 2 \mathrm{NO}(g)+\mathrm{Cl}_{2}(g) \rightleftharpoons 2 \mathrm{NOCl}(g) $$ for which \(\Delta H^{\circ}=-77.07 \mathrm{~kJ} .\) How will the amount of \(\mathrm{Cl}_{2}\) at equilibrium be affected by the following changes? (a) Removing \(\mathrm{NO}(g)\) (b) Adding \(\mathrm{NOCl}(g)\) (c) Raising the temperature (d) Decreasing the volume of the container at constant temperature

Short answer

(a) Amount of Cl2 will increase. (b) Amount of Cl2 will increase. (c) Amount of Cl2 will increase. (d) Amount of Cl2 will remain unchanged.

Step-by-step solution

- Applying Le Chatelier's Principle to Removal of NO

Le Chatelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. Removing NO(g) from the system reduces the concentration of NO on the reactant side, which will cause the equilibrium to shift to the left to produce more NO, thus the amount of Cl2 at equilibrium will increase.

- Applying Le Chatelier's Principle to Addition of NOCl

Adding NOCl(g) to the system increases its concentration on the product side. According to Le Chatelier's Principle, the equilibrium will shift to the left to reduce the concentration of NOCl by producing more reactants. Consequently, the amount of Cl2 at equilibrium will increase.

- Applying Le Chatelier's Principle to Raising the Temperature

Increasing the temperature of an exothermic reaction (one that releases heat) will cause the equilibrium to shift to the left, as the system tries to absorb the excess heat by favoring the endothermic direction. Since the reaction in the forward direction is exothermic, raising the temperature will result in an increase in the amount of Cl2 at equilibrium.

- Applying Le Chatelier's Principle to Decreasing the Volume

Decreasing the volume of the container at constant temperature increases the pressure of the system. The equilibrium will shift to the side with fewer moles of gas to decrease the pressure. Since the number of moles of gaseous products is the same as the reactants (2 moles of NO and 1 mole of Cl2 produce 2 moles of NOCl), there will be no shift in equilibrium and the amount of Cl2 at equilibrium will remain unchanged.

Key concepts

Chemical Equilibrium

Chemical equilibrium is a critical concept in chemistry which occurs when the rate of the forward reaction equals the rate of the reverse reaction in a chemical process, resulting in no net change in the concentrations of reactants and products over time. It's like a tug-of-war where both teams are equally strong, and the rope doesn't move. In a physical sense, the reactions are still occurring, but they counterbalance each other.

When a system is at equilibrium and experiences a change, be it in concentration, temperature, or pressure, Le Chatelier's Principle provides a predictive tool for understanding how the equilibrium will respond. Think of it as an attempt by the reaction to maintain balance or 'fight back' against the change. In the given exercise, we saw different scenarios where the equilibrium was disturbed, and how it sought to restore balance—such as by increasing the production of reactants when products were added or when NO was removed from the system.

Exothermic Reaction

An exothermic reaction is one that releases heat as it proceeds, resulting in a rise in temperature of the surroundings. It's like a cozy fireplace that warms the room on a chilly evening. In our case, the reaction converting Nitrogen monoxide (NO) and Chlorine (Cl2) to Nitrosyl chloride (NOCl) releases energy to the surroundings, indicated by a negative value of ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline (ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline (ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline �H^∘ = -77.07 kJewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline ewline �). When we talk about such a reaction in the context of Le Chatelier's Principle, increasing the temperature will shift the equilibrium in the direction that absorbs heat, which would be the reverse or endothermic direction in exothermic reactions. Hence, the exercise illustrated how raising the temperature causes the equilibrium to shift towards the reactants, increasing the amount of Cl2.

Reaction Quotient

The reaction quotient (Q) helps us determine the direction in which a reaction mixture will shift to reach equilibrium. To paint a clearer picture, Q is calculated using the same formula as the equilibrium constant (K), but with the initial concentrations or partial pressures of the reactants and products instead of the equilibrium concentrations. In essence, it's like taking a snapshot of where the system is headed at any point before equilibrium is reached. If the reaction quotient is less than K, the reaction moves forward to reach equilibrium. Conversely, if Q is greater than K, the reaction proceeds in the reverse direction. This concept wasn't directly examined in our previous exercise, but it lays the groundwork for understanding how changes in concentrations, temperature, or pressure can drive the system either towards or away from equilibrium.