The problem you're dealing with involves a chemical equilibrium for the exothermic reaction between sulfur dioxide (SO₂) and oxygen (O₂) to form sulfur trioxide (SO₃):

$2 \text{SO}_2 (g) + \text{O}_2 (g) \rightleftharpoons 2 \text{SO}_3 (g)$

Le Châtelier's principle tells us how the system at equilibrium will respond to changes in temperature and pressure.

1. Lowering the Temperature:

• Since the reaction is exothermic, heat can be considered a "product." Lowering the temperature favors the exothermic direction, shifting the equilibrium to the right (towards more SO₃).
• As a result, the pressure of SO₃ will increase.

2. Raising the Temperature:

• Raising the temperature will favor the endothermic direction, shifting the equilibrium to the left (towards more SO₂ and O₂).
• As a result, the pressure of SO₂ will increase.

Final Answers for the Table:

1. Lowering the Temperature:

   • The pressure of SO₃ will increase.
   • The equilibrium shifts to the right.

2. Raising the Temperature:

   • The pressure of SO₂ will increase.
   • The equilibrium shifts to the left.

Would you like further explanation on any part of this?

Here are five related questions you might consider:

1. How does increasing the concentration of SO₂ affect the equilibrium position?
2. What would happen if the volume of the reaction vessel is increased?
3. How does a catalyst affect the equilibrium position of this reaction?
4. How would the system respond if additional O₂ is introduced?
5. What are the real-life implications of controlling temperature in the production of SO₃?

Tip:

Le Châtelier's principle is a great tool to predict the direction of equilibrium shifts, but it does not give information about the rate at which equilibrium is re-established.