Question

Consider the following equilibrium reaction in a closed container: $$ \mathrm{CaCO}_{3}(s) \rightleftarrows \mathrm{CaO}(s)+\mathrm{CO}_{2}(g) $$ What will happen if (a) the volume is increased, (b) some \(\mathrm{CaO}\) is added to the mixture, \((\mathrm{c})\) some \(\mathrm{CaCO}_{3}\) is removed, \(\left(\right.\) d) some \(\mathrm{CO}_{2}\) is added to the mixture, (e) a few drops of an \(\mathrm{NaOH}\) solution are added to the mixture, (f) a few drops of an HCl solution are added to the mixture (ignore the reaction between \(\mathrm{CO}_{2}\) and water), (g) temperature is increased?

Short answer

(a) Right, (b) No change, (c) No change, (d) Left, (e) Right, (f) No change, (g) Right

Step-by-step solution

Understanding the Reaction

The reaction given is \( \text{CaCO}_3 (s) \rightleftarrows \text{CaO} (s) + \text{CO}_2 (g) \), an equilibrium involving solids and gas. Le Châtelier's principle will help in predicting the shift of equilibrium when changes are introduced. Remember that changes to solids don't affect equilibrium significantly, because their activities are constant.

Effect of Increasing Volume

Increasing the volume of the container will decrease the partial pressure of \( \text{CO}_2 \) (the only gaseous species). Le Châtelier's principle suggests that the system will shift towards the side with more gas molecules to counteract this change. Thus, the equilibrium will shift to the right, favoring the production of more \( \text{CO}_2 \).

Effect of Adding \( \text{CaO} \)

Adding more \( \text{CaO} \) does not affect the equilibrium position because it is a solid. Changes in the amount of solid do not influence the equilibrium of the reaction because their activities are constant in the reaction quotient.

Effect of Removing \( \text{CaCO}_3 \)

Similarly, removing \( \text{CaCO}_3 \) will not affect the equilibrium position for the same reason: solids have a constant activity and do not appear in the equilibrium expression.

Effect of Adding \( \text{CO}_2 \)

Adding \( \text{CO}_2 \) increases its partial pressure, causing the equilibrium to shift to the left by Le Châtelier's principle, thus favoring the production of \( \text{CaCO}_3 \) to reduce \( \text{CO}_2 \) partial pressure.

Effect of Adding \( \text{NaOH} \)

Adding \( \text{NaOH} \) solution introduces \( \text{OH}^- \) ions, which react with \( \text{CO}_2 \) to form \( \text{NaHCO}_3 \). This removal of \( \text{CO}_2 \) shifts the equilibrium to the right, according to Le Châtelier's principle.

Effect of Adding \( \text{HCl} \)

Adding \( \text{HCl} \) does not directly affect the equilibrium because \( \text{HCl} \) does not react with any of the components in this particular equilibrium system. Thus, there is effectively no change in the equilibrium position.

Effect of Increasing Temperature

For this reaction, the decomposition of \( \text{CaCO}_3 \) into \( \text{CaO} \) and \( \text{CO}_2 \) is an endothermic process. Increasing temperature will favor the forward reaction (endothermic direction), thus shifting the equilibrium to the right to absorb the added heat.

Key concepts

Equilibrium Reaction

In a chemical equilibrium reaction, the reactants and products exist in a dynamic balance where the rate of the forward reaction equals the rate of the reverse reaction. For the reaction \( \text{CaCO}_{3}(s) \rightleftarrows \text{CaO}(s) + \text{CO}_{2}(g) \), it involves both solids and a gas in a closed system. The concept of chemical equilibrium is foundational in understanding how the concentrations of reactants and products stabilize over time. It's important to note that the activities of solids are considered constant and do not show up in the equilibrium constant expression: \( K = [\text{CO}_2] \). This is why changes in solids like \( \text{CaO} \) and \( \text{CaCO}_{3} \) generally don't affect the position of equilibrium.

Effect of Volume Change

In a gaseous equilibrium, changing the volume of the container affects the pressure of the gases involved. When you increase the volume of the container for the equilibrium reaction \( \text{CaCO}_{3}(s) \rightleftarrows \text{CaO}(s) + \text{CO}_{2}(g) \), this decreases the partial pressure of \( \text{CO}_{2} \), the gaseous product. According to Le Châtelier's Principle, the equilibrium will shift to counteract the decrease in pressure. Thus, it will shift to the right to produce more \( \text{CO}_{2} \), increasing the total number of gas molecules. This is because the system aims to restore the pressure to what it was before the volume was increased.

Effects of Adding Reactants or Products

Adding or removing substances in an equilibrium reaction causes shifts in the equilibrium position because these changes affect the system's stress. For example, adding more \( \text{CO}_2 \) will increase its partial pressure, causing the equilibrium to shift to the left, producing more \( \text{CaCO}_{3} \) to reduce the \( \text{CO}_2 \) concentration. Conversely, adding \( \text{NaOH} \) will lead to a reaction with \( \text{CO}_2 \) to form \( \text{NaHCO}_{3} \), effectively removing \( \text{CO}_2 \) from the gaseous phase and shifting the equilibrium to the right to produce more \( \text{CO}_2 \). Interestingly, changes in pure solids like \( \text{CaO} \) and \( \text{CaCO}_{3} \) do not shift the equilibrium, as they do not appear in the equilibrium expression or change in concentration.

Impact of Temperature on Equilibrium

Temperature can significantly influence the direction of an equilibrium reaction. If a reaction is endothermic, it absorbs heat, while an exothermic reaction releases heat. In the case of \( \text{CaCO}_{3}(s) \rightleftarrows \text{CaO}(s) + \text{CO}_{2}(g) \), the reaction is endothermic, meaning it absorbs heat to break down \( \text{CaCO}_{3} \) into \( \text{CaO} \) and \( \text{CO}_{2} \). By increasing the temperature, more energy is available, favoring the forward reaction, which leads to an increase in the production of \( \text{CO}_{2} \). Hence, the equilibrium shifts to the right to absorb the additional heat provided, following Le Châtelier's Principle.