Question

Consider the reaction \(\mathrm{CaCO}_{3}(\mathrm{~s}) \rightleftharpoons \mathrm{CaO}(\mathrm{s})+\) \(\mathrm{CO}_{2}(\mathrm{~g})\) in a closed container at equilibrium. At a fixed temperature what will be the effect of adding more \(\mathrm{CaCO}_{3}\) on the equilibrium concentration of \(\mathrm{CO}_{2} ?\) (1) it increases (2) it decreases (3) it remains same (4) cannot be predicted unless the values of \(K_{p}\) is known

Short answer

It remains the same.

Step-by-step solution

Write the Equilibrium Expression

For the reaction \[\mathrm{CaCO}_{3}(\mathrm{~s}) \rightleftharpoons \mathrm{CaO}(\mathrm{s}) + \mathrm{CO}_{2}(\mathrm{~g})\]write the equilibrium expression. Since \(\mathrm{CaCO}_{3}\) and \(\mathrm{CaO}\) are solids, their concentrations are not included in the equilibrium expression. The equilibrium expression is presented in terms of the partial pressure of \(\mathrm{CO}_{2}\), denoted by \(P_{\mathrm{CO}_{2}}\):\[K_{p} = P_{\mathrm{CO}_{2}}\]

Analyze the Effect of Adding More Solid \(\mathrm{CaCO}_{3}\)

When more solid \(\mathrm{CaCO}_{3}\) is added, it does not change the concentration of the solid in the context of equilibrium. Therefore, adding more solid \(\mathrm{CaCO}_{3}\) does not affect the position of the equilibrium because the concentrations of pure solids are constants and do not appear in the equilibrium expression.

Determine the Effect on Equilibrium Concentration of \(\mathrm{CO}_{2}\)

Since the addition of solid \(\mathrm{CaCO}_{3}\) does not alter the equilibrium position, the equilibrium concentration of \(\mathrm{CO}_{2}\) remains unchanged. Thus, the equilibrium concentration of \(\mathrm{CO}_{2}\) is unaffected by adding more \(\mathrm{CaCO}_{3}\).

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle helps us understand how a system at equilibrium responds to disturbances. When a change occurs in the system, the equilibrium shifts to counteract that change, striving to restore balance. This principle is crucial in chemical equilibrium studies.
For example, if we add more reactants or remove products, the system shifts to create more products, and vice versa. This principle is all about minimizing the effects of changes imposed on the chemical equilibrium.

Equilibrium Expression

In chemical reactions, the equilibrium expression represents a ratio of the concentrations of products to reactants at equilibrium. This is simplified by using partial pressures for gases. For the given reaction: \(\text{CaCO}_{3}(\text{s}) \rightleftharpoons \text{CaO}(\text{s}) + \text{CO}_{2}(\text{g})\), the equilibrium expression only considers gases because the concentrations of solids are constants.
Thus, the equilibrium constant, \(K_p\), is expressed in terms of the partial pressure of \(\text{CO}_{2}\). The balanced equilibrium equation is given by: \[ \text{K}_{\text{p}} = \text{P}_{\text{CO}_{2}}\] By consistently relating partial pressures to the equilibrium expression, predictions about system responses to changes can be made using Le Chatelier's Principle.

Partial Pressure

Partial pressure refers to the pressure exerted by a single gas component in a mixture of gases. In equilibrium studies, partial pressure is important because it reflects the tendency of a gas to escape or be absorbed. For example, in the reaction: \[ \text{CaCO}_{3}(\text{s}) \rightleftharpoons \text{CaO}(\text{s}) + \text{CO}_{2}(\text{g})\], the partial pressure of \(\text{CO}_{2}\) (\text{P}_{\text{CO}_{2}}) indicates the amount of carbon dioxide gas present. Changes in the partial pressure directly affect the equilibrium state, as described by the equilibrium expression.