Question

Explain how you would regulate the pressure to favor the products in the following equilibrium system. $$ \mathrm{MgCO}_{3}(\mathrm{s}) \rightleftharpoons \mathrm{MgO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g}) $$

Short answer

To favor the formation of MgO and CO2 in the given equilibrium system, we should decrease the pressure on the system. This is because there are more moles of gas on the product side (1 mole of CO2). According to Le Chatelier's principle, reducing the pressure shifts the equilibrium position in the direction of the side with more moles of gas, thus favoring the production of magnesium oxide and carbon dioxide.

Step-by-step solution

Identify the equilibrium system

The equilibrium system is given by the following equation: \[ \mathrm{MgCO}_{3}(\mathrm{s}) \rightleftharpoons \mathrm{MgO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g}) \] In this equilibrium system, magnesium carbonate (MgCO3) decomposes into magnesium oxide (MgO) and carbon dioxide (CO2). Here, MgCO3 and MgO are solid, and CO2 is a gas.

Apply Le Chatelier's principle

According to Le Chatelier's principle, when the pressure on an equilibrium system is increased, the system will shift its position in favor of the side with fewer moles of gas. When the pressure is decreased, it will shift in favor of the side with more moles of gas. For our given system, On the reactant side, there is 1 mole of solid (MgCO3), and on the product side, there is 1 mole of solid (MgO) and 1 mole of gas (CO2).

Regulate the pressure to favor the products

As there are more moles of gas on the product side of the equation (1 mole of CO2), we need to decrease the pressure to favor the formation of products (MgO and CO2). By reducing the pressure, the equilibrium position shifts in the direction of the side with more moles of gas, as per Le Chatelier's principle. Therefore, the production of magnesium oxide and carbon dioxide will be favored. In summary, to favor the formation of MgO and CO2 in the given equilibrium system, we should decrease the pressure on the system.

Key concepts

Equilibrium Systems

An equilibrium system in chemistry refers to a state where the rate of the forward reaction equals the rate of the reverse reaction. This means that the concentrations of reactants and products remain constant over time. Equilibrium is dynamic, meaning that the reactions continue to occur, but there is no net change in their concentrations. In the given chemical equation \[\mathrm{MgCO}_{3} (s) \rightleftharpoons \mathrm{MgO} (s) + \mathrm{CO}_{2} (g)\]MgCO₃ is decomposing into MgO and CO₂. The system reaches equilibrium when the rate of decomposition equals the rate of formation of magnesium carbonate from magnesium oxide and carbon dioxide. Understanding equilibrium systems allows chemists to predict the effects of changing conditions on the concentrations of substances within a reaction.

Pressure Regulation

Pressure regulation in equilibrium systems is an important concept governed by Le Chatelier's Principle. This principle states that if a dynamic equilibrium system is disturbed by changing the conditions, such as pressure, the system will adjust itself to partially counteract the disturbance and restore a new equilibrium.For reactions involving gases, changes in pressure can significantly alter the equilibrium position. In the reaction \( \mathrm{MgCO}_{3} \rightleftharpoons \mathrm{MgO} + \mathrm{CO}_{2} \), we can manipulate the pressure to influence the production of gases.

• Increasing pressure: The system shifts toward the side with fewer moles of gas to reduce the added pressure.
• Decreasing pressure: The system shifts toward the side with more moles of gas, favoring gas production.

In this case, by decreasing the pressure, the equilibrium will shift to the right, favoring the formation of \( \mathrm{MgO} \) and \( \mathrm{CO}_{2} \), thus increasing the production of products.

Chemical Decomposition

Chemical decomposition is the process where a single compound breaks down into two or more simpler substances. It is an essential type of reaction, especially in reactions like the decomposition of magnesium carbonate.In the given reaction, \( \mathrm{MgCO}_{3} \) decomposes into \( \mathrm{MgO} \) and \( \mathrm{CO}_{2} \). This type of reaction generally requires energy to break chemical bonds, often supplied in the form of heat.Decomposition reactions are crucial for understanding how complex substances can yield simpler products. This understanding is key in areas such as material processing, combustion, and environmental science, where decomposition reactions often play a central role.

Mole Concept

The mole concept is a fundamental idea in chemistry used to quantify the amount of substance involved in a reaction. A mole is defined as the amount of substance containing the same number of entities, such as atoms or molecules, as there are in 12 grams of carbon-12. This number is known as Avogadro's number: approximately \( 6.022 \times 10^{23} \) entities.In equilibrium reactions, understanding the mole concept helps to predict the direction of the shift in equilibrium when changes occur. For example, knowing that there is 1 mole of \( \mathrm{CO}_{2} \) on the product side in the reaction \( \mathrm{MgCO}_{3} \rightarrow \mathrm{MgO} + \mathrm{CO}_{2} \) helps in applying Le Chatelier's Principle effectively.By calculating and comparing moles on each side, chemists can predict how changes in pressure, volume, or temperature will affect an equilibrium state. This is crucial for reactions where gases are involved, as their volumes and pressures are directly related to the number of moles present.