Question

Given this equilibrium, predict the direction of shift for each stress listed. \(\mathrm{CO}_{2}(\mathrm{~g})+\mathrm{C}(\mathrm{s})+171 \mathrm{~kJ} \rightleftarrows 2 \mathrm{CO}(\mathrm{g})\) a) addition of \(\mathrm{CO}\) b) increased pressure c) addition of a catalyst

Short answer

a) Left, b) Left, c) No shift.

Step-by-step solution

Identify the Type of Reaction

This is an endothermic reaction, as indicated by the positive enthalpy change (+171 kJ), which means heat is absorbed by the system. The equilibrium shifts according to changes in concentration, pressure, and temperature to counteract these changes, as stated by Le Chatelier's Principle.

Analyze Effect of Adding CO

According to Le Chatelier's Principle, if a concentration of a product is increased, the equilibrium will shift in the direction that consumes the added product. Therefore, when \(\mathrm{CO}\) is added to the system, the equilibrium will shift towards the left to form \(\mathrm{CO}_2(\mathrm{~g})+\mathrm{C}(\mathrm{s})\).

Examine Impact of Increased Pressure

When pressure is increased, the equilibrium will shift towards the side with fewer moles of gas to reduce pressure. In the given reaction, there are more gas moles on the right side (2 moles of CO) than on the left side (1 mole of \(\mathrm{CO}_2\)). Thus, an increase in pressure will shift the equilibrium to the left.

Assess the Role of a Catalyst

The addition of a catalyst speeds up the rate of both the forward and reverse reactions equally without affecting the position of the equilibrium. Therefore, the equilibrium position will not shift with the addition of a catalyst.

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle helps us predict how changes in a system's conditions will affect chemical equilibrium. Imagine a seesaw balancing perfectly. If someone adds weight to one side, you need to shift to balance it again. This is similar to how equilibrium shifts.

When something in the system changes—like concentration, temperature, or pressure—the equilibrium will adjust to counteract that change and reach a new balance. So, if you add a product like \(\mathrm{CO}\) in our equilibrium reaction, the system will try to "remove" that extra product by shifting the balance to the left side, forming more \(\mathrm{CO}_2(\mathrm{~g})+\mathrm{C}(\mathrm{s})\).

• An increase in reactants will shift equilibrium towards products.
• An increase in products will shift equilibrium towards reactants.

This balancing act maintains equilibrium when changes occur, like keeping a seesaw level.

Endothermic Reaction

Endothermic reactions occur when a system absorbs energy, typically in the form of heat. Think of it like the reaction needs extra heat to proceed, much like needing more fuel to power a machine.

In our provided reaction, \(\mathrm{CO}_{2}(\mathrm{~g})+\mathrm{C}(\mathrm{s})+171 \mathrm{~kJ} \rightleftarrows 2 \mathrm{CO}(\mathrm{g})\), the positive energy term (+171 kJ) indicates it requires heat. This absorbed heat becomes a reactant in the reaction.

• Adding heat shifts equilibrium to the right, favoring the formation of \(\mathrm{CO}\).
• Removing heat will shift it to the left, opposing the endothermic process.

This concept aligns with Le Chatelier's Principle, as the system adapts to change by either taking in or releasing heat to maintain balance.

Catalysts in Equilibrium

Catalysts are like chemical speed boosters. They make reactions happen faster, but don't affect where equilibrium sits. Think of them as friends who help you meet halfway faster but don't change the meeting point itself.

In our chemical system, adding a catalyst does not shift equilibrium. Instead, it equally speeds up both the forward and reverse reactions, allowing the system to reach equilibrium more quickly.

• They lower the activation energy, making it easier for reactions to occur.
• The position of equilibrium remains unchanged.

Thus, while a catalyst makes the reaction more efficient, the ratio of products to reactants in equilibrium stays the same.

Gas Pressure and Equilibrium

Pressure changes can sway the direction of gaseous reactions. In gases, volume and pressure are tightly linked, so changing the number of gas moles affects equilibrium.

In our specific reaction, increasing pressure shifts equilibrium towards the side with fewer gas moles. As \(\mathrm{CO}_2(\mathrm{~g}) \) has 1 mole on the left and \(2\,\mathrm{CO}(\mathrm{g})\) has 2 moles on the right, increasing pressure shifts the system to the left.

• Decreasing pressure shifts to the side with more gas moles.
• Increasing pressure shifts to the side with fewer gas moles.

This principle ensures systems are more comfortable—like adjusting seating in a crowded room by moving people to emptier spaces.