Consider the endothermic reaction: $$ \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{~g})+\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{l}_{2}(\mathrm{~g}) $$ If you were trying to maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{l}_{2}\) produced, which tactic might you try? Assume that the reaction mixture reaches equilibrium. a. decreasing the reaction volume b. removing \(1_{2}\) from the reaction mixture c. raising the reaction temperature d. adding \(\mathrm{C}_{2} \mathrm{H}_{4}\) to the reaction mixture

In endothermic reactions, energy in the form of heat is absorbed from the surroundings to break the bonds of the reactants, causing the reaction container to feel cold. These reactions require heat to proceed, and their progress can be influenced significantly by temperature changes.

Example: The production of ethylene iodide from ethylene and iodine is endothermic. When heat is added to the reaction system, the reaction absorbs it, and more products are formed as the reaction moves forward to consume the additional energy.

Chemical equilibrium refers to a state in a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction, leading to constant concentrations of reactants and products. This balance can be disrupted, though, by changes in concentration, temperature, or pressure.

Within a closed system at equilibrium, the reaction can appear to have stopped, but in reality, both the forward and reverse reactions are still occurring at the same rate. It's a dynamic state, not a static one.

The temperature at which a reaction is performed can play a critical role in determining the position of equilibrium. In an endothermic process, an increase in temperature provides energy for the reaction to proceed, thus favoring the formation of products.

Practical Application: By raising the temperature of the reaction between ethylene and iodine, chemists can drive the equilibrium toward more ethylene iodide formation, using temperature to control the yield of the desired product.

An equilibrium shift occurs when a change in conditions causes a reaction at equilibrium to favor either the forward or reverse process in an attempt to counteract the imposed change.

This behavior is described by Le Chatelier's principle, which states that if a system at equilibrium is disturbed, it will adjust to diminish the change. For instance, removing one of the reactants, like iodine in the given reaction, would typically cause a shift to the left as the system attempts to replace the iodine, but this would not be the correct approach to maximize the product. Instead, increasing temperature for endothermic reactions, like adding heat to our ethylene and iodine reaction, shifts the equilibrium to the right, thereby increasing product formation.