Question

State Le Chatlier's principle and apply it to the following equilibria: (i) \(2 \mathrm{NO}_{(\mathrm{g})} \rightleftarrows \mathrm{N}_{2(\mathrm{~g})}+\mathrm{O}_{2(\mathrm{~g})}+\) Heat (ii) \(\mathrm{N}_{2} \mathrm{O}_{4(\mathrm{~g})} \rightleftarrows 2 \mathrm{NO}_{2(\mathrm{~g})}-\) Heat

Short answer

2. How does the equilibrium of reaction (i) shift when the concentration of NO is increased or decreased? 3. How does the equilibrium of reaction (i) shift when the temperature is increased or decreased? 4. How does the equilibrium of reaction (ii) shift when the concentration of N2O4 is increased or decreased? 5. How does the equilibrium of reaction (ii) shift when the temperature is increased or decreased?

Step-by-step solution

State Le Chatelier's Principle

Le Chatelier's Principle states that if a chemical system at equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium will shift in a direction that minimizes the impact of that change. In other words, the system will adjust to counteract the disturbance and restore a new equilibrium.

Equilibrium (i): \(2 \mathrm{NO}_{(\mathrm{g})} \rightleftarrows \mathrm{N}_{2(\mathrm{~g})}+\mathrm{O}_{2(\mathrm{~g})}+\) Heat

In order to analyze this equilibrium, let's categorize the changes that could occur in the system: 1. Increasing the concentration of \(\mathrm{NO}\): The equilibrium will shift to the right to consume the excess \(\mathrm{NO}\). 2. Decreasing the concentration of \(\mathrm{NO}\): The equilibrium will shift to the left to produce more \(\mathrm{NO}\). 3. Increasing the concentration of \(\mathrm{N_{2}}\) or \(\mathrm{O_{2}}\): The equilibrium will shift to the left to consume the excess \(\mathrm{N_{2}}\) or \(\mathrm{O_{2}}\). 4. Decreasing the concentration of \(\mathrm{N_{2}}\) or \(\mathrm{O_{2}}\): The equilibrium will shift to the right to produce more \(\mathrm{N_{2}}\) or \(\mathrm{O_{2}}\). 5. Increasing temperature (adding heat): Since the reaction is endothermic (+ Heat), the position will shift to the right to absorb the heat, resulting in an increase in product formation. 6. Decreasing temperature (removing heat): The position will shift to the left to produce heat, consuming the products and increasing the reactants.

Equilibrium (ii): \(\mathrm{N_{2}O_{4}}_{(\mathrm{g})} \rightleftarrows 2 \mathrm{NO}_{2(\mathrm{g})}-\) Heat

Similar to the previous equilibrium, we will analyze the changes that could occur in the system: 1. Increasing the concentration of \(\mathrm{N_{2}O_{4}}\): The equilibrium will shift to the right to consume the excess \(\mathrm{N_{2}O_{4}}\). 2. Decreasing the concentration of \(\mathrm{N_{2}O_{4}}\): The equilibrium will shift to the left to produce more \(\mathrm{N_{2}O_{4}}\). 3. Increasing the concentration of \(\mathrm{NO_{2}}\): The equilibrium will shift to the left to consume the excess \(\mathrm{NO_{2}}\). 4. Decreasing the concentration of \(\mathrm{NO_{2}}\): The equilibrium will shift to the right to produce more \(\mathrm{NO_{2}}\). 5. Increasing temperature (adding heat): Since the reaction is exothermic (- Heat), the position will shift to the left to offset the added heat, resulting in an increase in reactant formation. 6. Decreasing temperature (removing heat): The position will shift to the right to replace the lost heat, consuming the reactants and increasing the products.

Key concepts

Chemical Equilibrium

Chemical equilibrium is a state in any reversible chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of all reactants and products remain constant over time, but not necessarily equal. This equilibrium can be visually represented by a double-headed arrow in the reaction equation.

This balance, however, does not mean the reactions have stopped; rather, the chemical processes continue to occur in both directions at an equal pace. Hence, there is a dynamic exchange at the molecular level, but macroscopically we observe a stable system. A vital concept to understand in chemical equilibrium is that it's a balance of rates, not amounts..

Concentration Effect

The concentration effect, according to Le Chatelier's principle, has direct implications on chemical equilibrium. When the concentration of a reactant in a balanced system increases, the system seeks equilibrium by shifting to consume the added substance, thus favoring the forward reaction. Conversely, if a product's concentration increases, the system will shift to favor the reverse reaction to reduce this concentration.

These shifts are the system's attempt to maintain equilibrium by adjusting concentrations. In educational exercises, students can predict the outcome of altering concentrations: adding reactants leads to more products, and adding products leads to more reactants.

Temperature Effect on Equilibrium

Temperature changes can profoundly influence chemical equilibrium. An increase in temperature adds energy to the system, affecting endothermic and exothermic reactions differently. An endothermic reaction, where heat is absorbed, will shift towards the products as it consumes the extra heat. In contrast, an exothermic reaction, where heat is released, will shift towards the reactants to absorb the excess heat energy.

On cooling, the system will trend in the opposite direction. For endothermic reactions, the equilibrium will shift towards the reactants to produce more heat, while for exothermic reactions, it will shift towards the products to release heat. This concept allows students to predict how temperature alterations can change reaction yields in both directions.

Pressure Effect on Equilibrium

Pressure changes can also affect the equilibrium of reactions involving gases, as per Le Chatelier's principle. An increase in pressure typically causes the equilibrium position to shift towards the side with fewer moles of gas, as this reduces the volume and therefore the pressure. Conversely, a decrease in pressure favors the side with more moles of gas.

This rule, however, only applies when gases are involved and there’s a difference in the number of moles between reactants and products. By understanding this, students can solve problems related to equilibrium by considering the number of gas molecules on each side of a balanced reaction equation.