The equilibrium reaction given is: \[ \text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g) \] ### Correct Option: C **Step-by-Step Explanation:** 1. **Understanding the Reaction:** - The reaction involves the conversion of dinitrogen tetroxide (N₂O₄) into nitrogen dioxide (NO₂). - The balanced equation shows that 1 mole of N₂O₄ produces 2 moles of NO₂. 2. **Le Chatelier's Principle:** - This principle states that if a system at equilibrium is subjected to a change in pressure, temperature, or concentration, the system will adjust to counteract that change and restore a new equilibrium. - In this case, we are focusing on the effect of an increase in total pressure. 3. **Analyzing Moles of Gas:** - In the reaction, we have: - 1 mole of N₂O₄ (reactant) - 2 moles of NO₂ (product) - Therefore, the total number of moles of gas on the left side (reactants) is 1, and on the right side (products) is 2. 4. **Effect of Increasing Pressure:** - When the total pressure of the system is increased, the equilibrium will shift towards the side with fewer moles of gas to reduce the pressure. - In this case, since there are fewer moles of gas on the left side (1 mole of N₂O₄) compared to the right side (2 moles of NO₂), the equilibrium will shift to the left, favoring the formation of N₂O₄. 5. **Conclusion:** - Therefore, an increase in total pressure will produce more N₂O₄ in the mixture, making option D the correct answer. ### Why Other Options Are Incorrect: - **Option A: Produce more of NO₂(g) in the mixture** - This option is incorrect because increasing the pressure favors the side with fewer moles of gas, which is the reactant side (N₂O₄). Thus, NO₂ will not be produced in greater amounts. - **Option B: Convert all of N₂O₄ to NO₂(g)** - This option is also incorrect. While the reaction can proceed to produce NO₂, an increase in pressure does not guarantee that all N₂O₄ will convert to NO₂. The equilibrium will shift to favor the formation of N₂O₄, not completely convert it. - **Option C: Have no effect on the concentrations of NO₂(g) and N₂O₄(g)** - This option is incorrect because an increase in pressure does affect the equilibrium position. It will shift towards the side with fewer moles of gas, which is the reactant side (N₂O₄). ### Summary of Key Points: - Le Chatelier's Principle helps predict how a system at equilibrium responds to changes in pressure. - Increasing pressure shifts the equilibrium towards the side with fewer moles of gas. - In the reaction N₂O₄ ↔ 2NO₂, the left side has fewer moles (1 mole) compared to the right side (2 moles). - Therefore, increasing pressure will favor the formation of N₂O₄, not NO₂. This understanding is crucial for predicting the behavior of gas-phase reactions under varying conditions.