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

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.

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.

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.