Question

Consider the exothermic reaction: $$ \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{Q}_{2}(g) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}(\mathrm{~g}) $$ If you were trying to maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) produced, which tactic might you try? Assume that the reaction mixture reaches equllubrium. a. increasing the reaction volume b. removing \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) from the reaction mixture as it forms c. lowering the reaction temperature d. adding \(\mathrm{Cl}_{2}\)

Short answer

To maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) produced, the best tactic is to remove \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) from the reaction mixture as it forms, shifting the equilibrium towards more product formation.

Step-by-step solution

Identify the Nature of the Reaction

Determine whether the reaction is exothermic or endothermic. In this case, it is given that the reaction is exothermic, meaning it releases heat.

Apply Le Chatelier's Principle

Use Le Chatelier's Principle to predict how the system at equilibrium will respond to changes. This principle states that if a change is applied to a system at equilibrium, the system adjusts in a way that minimizes the change.

Analyze Each Tactic

Consider each tactic and predict its effect on the reaction equilibrium:a. Increasing the reaction volume would generally shift the equilibrium towards the side with more gaseous molecules, which is not applicable here because the number of moles of gases is equal on both sides of the reaction.b. Removing \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) as it forms would shift the equilibrium to the right (product side) to replace the removed product.c. Lowering the temperature for an exothermic reaction would shift the equilibrium to the right to produce more heat (product side).d. Adding \(\mathrm{Cl}_{2}\) would shift the equilibrium to the right because of the increase in the concentration of a reactant, resulting in more product formation.

Choose the Best Tactic

Select the tactic that will maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) produced. Removing the product, \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\), will continuously shift the equilibrium towards the product side to replenish the removed compound, resulting in more product formation.

Key concepts

Chemical Equilibrium

Understanding chemical equilibrium is essential to manipulating any chemical reaction for desired outcomes. It refers to the state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of the reactants and products. However, this does not mean that the reactants and products are in equal amounts, rather that they exist in a fixed ratio.

When a system reaches equilibrium, it may seem like the reaction has stopped, but in reality, both the forward and reverse reactions are still occurring; they're simply in a state of balance. It's as if both teams in a tug-of-war are equally strong, and the rope does not move to either side. For the reaction given in the exercise, achieving the point where \(\mathrm{C}_2\mathrm{H}_4(g)+\mathrm{Cl}_2(g) \) and \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2(g) \) are produced and consumed at an equal rate is the goal to approach chemical equilibrium.

To manipulate the equilibrium in your favor, for instance, to produce more \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2\), you could change factors like concentration, temperature, or pressure of the system where principles such as Le Chatelier's come into play.

Exothermic Reaction

An exothermic reaction is a type of chemical reaction that releases energy, usually in the form of heat, to its surroundings. This contrasts with endothermic reactions, which absorb energy. Imagine exothermic reactions like a cozy campfire, emitting warmth into the cool night air.

In the context of the given exercise, the reaction of \(\mathrm{C}_2\mathrm{H}_4(g) + \mathrm{Cl}_2(g) \) producing \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2(g) \) and releasing energy is an exothermic process. Thus, any change to the system that removes heat (like lowering the temperature) will shift the equilibrium toward producing more heat; meaning more products will be formed as well to compensate for the heat loss according to Le Chatelier's Principle. This concept is pivotal when considering how to alter the reaction conditions to increase the yield of \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2\). For instance, increasing the temperature would have the opposite effect, favoring the reactants instead.

Equilibrium Shift

The term equilibrium shift refers to the change in the position of equilibrium when a system in equilibrium is subjected to an external change. According to Le Chatelier's Principle, the system tends to adjust to minimize the effect of this change to re-establish equilibrium.

For example, if you increase the concentration of a reactant or decrease the concentration of a product, the equilibrium will 'shift' to produce more products. Likewise, decreasing the concentration of a reactant or increasing the concentration of a product will 'shift' the equilibrium the other way, to favor reactants.

Applying this to our exercise, tactics such as removing the product \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2\) as it forms or adding more of a reactant like \(\mathrm{Cl}_2\), drive the reaction forward, effectively shifting the equilibrium to the right to produce more \(\mathrm{C}_2\mathrm{H}_4\mathrm{Cl}_2\). These approaches are examples of altering the equilibrium to optimize the yield of the desired chemical, which is crucial in industrial chemistry for maximizing production efficiency.