Question

The industrial process for producing carbon monoxide gas involves passing carbon dioxide over hot charcoal. $$ \mathrm{C}(s)+\mathrm{CO}_{2}(g)+\text { heat } \rightleftarrows 2 \mathrm{CO}(g) $$ Predict the direction of equilibrium shift for each of the following stresses: (a) increase \(\left[\mathrm{CO}_{2}\right]\) (b) decrease \(\left[\mathrm{CO}_{2}\right]\) (c) increase [CO] (d) decrease \([\mathrm{CO}]\) (e) increase temperature (f) decrease temperature (g) increase volume (h) decrease volume (i) add C powder (j) add Kr inert gas

Short answer

(a) Right, (b) Left, (c) Left, (d) Right, (e) Right, (f) Left, (g) Right, (h) Left, (i) No effect, (j) No effect.

Step-by-step solution

Identifying Reaction Type

The given reaction is an endothermic process, where carbon dioxide reacts with carbon to form carbon monoxide. The reaction can be represented as \( \mathrm{C}(s) + \mathrm{CO}_2(g) + \text{heat} \rightleftarrows 2\mathrm{CO}(g) \).

Analyzing Effect of Concentration Changes on Equilibrium (a)

When \( [\mathrm{CO}_2] \) is increased, the system will shift to the right to reduce \( [\mathrm{CO}_2] \), forming more \( \mathrm{CO} \).

Analyzing Effect of Concentration Changes on Equilibrium (b)

When \( [\mathrm{CO}_2] \) is decreased, the system will shift to the left to increase \( [\mathrm{CO}_2] \), leading to the formation of more reactants.

Analyzing Effect of Concentration Changes on Equilibrium (c)

Increasing \( [\mathrm{CO}] \) will cause the system to shift to the left to reduce \( [\mathrm{CO}] \), favoring the formation of \(\mathrm{CO}_2 \).

Analyzing Effect of Concentration Changes on Equilibrium (d)

Decreasing \( [\mathrm{CO}] \) will cause the system to shift to the right to increase \( [\mathrm{CO}] \), producing more \( \mathrm{CO} \).

Analyzing Effect of Temperature Changes on Equilibrium (e)

Increasing temperature in an endothermic reaction shifts the equilibrium to the right, favoring the formation of \( \mathrm{CO} \).

Analyzing Effect of Temperature Changes on Equilibrium (f)

Decreasing temperature in an endothermic reaction shifts the equilibrium to the left, favoring the formation of \( \mathrm{CO}_2 \).

Analyzing Effect of Volume Changes on Equilibrium (g)

Increasing volume will cause the system to shift towards the side with more gas moles to occupy the additional volume, which is the right side (produces \( \mathrm{CO} \)).

Analyzing Effect of Volume Changes on Equilibrium (h)

Decreasing volume will cause the system to shift towards the side with fewer gas moles, which is the left side (reactants).

Effect of Adding a Solid on Equilibrium (i)

Adding more solid \( \mathrm{C} \) does not affect equilibrium since the concentration of solids does not appear in the equilibrium expression.

Effect of Adding an Inert Gas on Equilibrium (j)

Adding \( \mathrm{Kr} \), an inert gas, at constant volume does not change the partial pressures of reacting gases, so equilibrium position remains unchanged.

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that predicts how a chemical system at equilibrium will respond to changes or "stresses" in conditions. As per this principle, if an external change is applied to a system, the equilibrium will shift in a direction that counteracts the imposed change. This helps maintain balance in the system.

When you increase the concentration of a reactant like \([\mathrm{CO}_2]\), for example, the system will try to use up the added reactant by shifting in the direction that produces more products, so it moves to the right.

Similarly, changes in temperature, pressure, and concentration can all cause shifts in equilibrium. This principle helps predict both the direction and extent of these shifts to maintain balance over time.

Endothermic Reaction

An endothermic reaction is a type of chemical reaction that absorbs energy from its surroundings. This energy is usually in the form of heat, leading these reactions to feel cold to the touch.

In the reaction \(\mathrm{C}(s) + \mathrm{CO}_2(g) + \text{heat} \rightleftarrows 2 \mathrm{CO}(g)\), heat acts as a reactant. When the temperature of an endothermic reaction is increased, the equilibrium shifts to the right to absorb the extra heat, forming more products in the process.

This is why increasing the temperature of the carbon monoxide production reaction causes more \(\mathrm{CO}\) to form. Conversely, lowering the temperature removes heat energy, shifting the equilibrium towards reactants.

Equilibrium Shift

The equilibrium shift refers to the movement of a chemical equilibrium in response to changes in conditions such as concentration, temperature, or pressure.

For example, increasing the concentration of \([\mathrm{CO}_2]\) causes a shift to the right. This means more \(\mathrm{CO}\) is created, moving towards the products. If the concentration of \([\mathrm{CO}\)] is increased, the reaction favors the reverse direction, moving towards the left, thereby forming more \(\mathrm{CO}_2\).

The equilibrium shift is a dynamic process that allows chemical systems to adapt and reach a new state of balance under changed external conditions. The direction of shifts can usually be predicted by Le Chatelier's principle.

Effect of Volume on Equilibrium

The effect of volume changes on equilibrium is another crucial aspect when understanding how gaseous reactions react to changes. According to Le Chatelier's Principle, if you increase the volume of a reaction chamber containing gases, the equilibrium will shift towards the direction that produces more gas molecules or takes up more space.

In the given reaction \(\mathrm{C}(s) + \mathrm{CO}_2(g) + \text{heat} \rightleftarrows 2 \mathrm{CO}(g)\), increasing the volume of the container would cause the equilibrium to shift to the right because this side has more moles of gas \((2 \mathrm{CO})\) compared to only one mole of \(\mathrm{CO}_2\) on the left.

Conversely, decreasing the volume causes the reaction to shift to the left side, reducing the number of gaseous molecules, thus favoring the formation of reactants.