Question

The following reaction is exothermic. $$ \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{Br}_{2}(g) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Br}_{2}(g) $$ Predict the effect (shift right, shift left, or no effect) of these changes. (a) increasing the reaction temperature (b) decreasing the reaction temperature

Short answer

Increasing the reaction temperature will shift the equilibrium to the left, while decreasing the temperature will shift it to the right.

Step-by-step solution

Understanding Le Chatelier's Principle

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. For an exothermic reaction, increasing temperature will shift the equilibrium to the left (to absorb the excess heat). Decreasing temperature will shift the equilibrium to the right (to release heat and raise the temperature).

Predict the Effect of Increasing Temperature

Increasing the temperature adds kinetic energy to the system which is equivalent to adding heat. Since the reaction is exothermic, the system will counteract this by favoring the reverse (endothermic) reaction that absorbs heat. Thus, the equilibrium will shift to the left.

Predict the Effect of Decreasing Temperature

Decreasing the temperature removes kinetic energy from the system, effectively removing heat. The system responds by shifting the equilibrium toward the exothermic direction to produce more heat. Therefore, the equilibrium will shift to the right.

Key concepts

Exothermic Reaction

An exothermic reaction is a type of chemical process that releases energy, usually in the form of heat, to its surroundings. This is a phenomenon commonly encountered in everyday life, such as when wood burns in a fireplace or when an instant heat pack is activated. In these reactions, the energy needed to break the bonds of the reactants is less than the energy released when new bonds form in the products.

The general equation for an exothermic reaction can be depicted as:
\[\[\begin{align*}\text{Reactants} \rightarrow \text{Products} + \text{Energy (heat)}\tag{1}\end{align*}\]\]
As energy is a product, the reaction inherently favors the formation of the products when energy is removed from the system. Conversely, adding energy to the system tends to be unfavorable to the forward reaction. In a classroom setting, these processes can be illustrated through experiments that involve temperature changes observable by touch or with a thermometer.

Chemical Equilibrium

Chemical equilibrium is the state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction. This results in the concentrations of the reactants and products remaining constant over time, not because the reactions have stopped, but because they are going on at exactly the same rate.

The mathematical expression for the equilibrium constant (\text{K}_{\text{eq}}) provides insight into the relative amounts of reactants and products at equilibrium:
\[\[\begin{align*}\text{K}_{\text{eq}} = \frac{\text{[Products]}}{\text{[Reactants]}}\tag{2}\end{align*}\]\]
While the value of \text{K}_{\text{eq}} does not inform us about the specifics of reaction rates, it is a vital clue to the position of equilibrium. A high \text{K}_{\text{eq}} value suggests a product-favored reaction, while a low value indicates a reactant-favored reaction. Understanding how equilibrium may shift in response to external changes is critical for students studying chemical reactions.

Equilibrium Shift

An equilibrium shift occurs when an external force disturbs a system that is in chemical equilibrium, causing it to favor either the forward or the reverse reaction in order to reestablish a new equilibrium state. Le Chatelier's Principle is key to predicting these shifts.

Changes that can cause a shift include alterations in concentration, pressure, and temperature. If a reactant or product is added or removed from the system, the equilibrium will shift to oppose this change – for example, adding more reactants typically causes a shift to the right, increasing the amount of products. A reduction in volume, which increases pressure, will shift the equilibrium toward the side with fewer moles of gas, and vice versa for an increase in volume.

Understanding these shifts is not just academic; they have real-world implications in industries like pharmaceuticals where yields of a desired chemical can be maximized by controlling equilibrium conditions.

Reaction Temperature Effect

The effect of reaction temperature on chemical equilibrium is profound and is directly linked to the exo- or endothermic nature of the reactions involved. For exothermic reactions, an increase in temperature supplies more heat to the system, which the reaction naturally releases. According to Le Chatelier’s Principle, the equilibrium will shift to the left, favoring the reverse, endothermic reaction to remove the added heat and restore equilibrium.

Conversely, when the temperature decreases, the system loses heat energy. To counteract this, the equilibrium shifts to the right, favoring the forward, exothermic reaction to produce more heat.

This temperature dependence is critical in industries where controlling reaction conditions is essential. For example, in chemical manufacturing, carefully managing the temperature can optimize the yield and efficiency of product formation.