Final Answer:
Adding more sulfur trioxide (SO3) would prevent the reaction from shifting towards it, as it would oppose the forward reaction and push the equilibrium back towards the reactants.
Step-by-step explanation:
Le Chatelier's principle states that a system at equilibrium will adjust to oppose any changes in its conditions. In this reversible reaction, adding more SO3 (product) would increase its concentration, driving the reaction backwards to decrease its concentration and maintain equilibrium. Therefore, adding more SO3 would prevent the reaction from shifting further towards forming it.
Other options might seem appealing but are incorrect:
Removing SO3: This would directly shift the equilibrium towards forming more SO3 to replenish the removed product.
Adding sulfur dioxide (SO2): This is a reactant, and adding more would drive the reaction forward, producing more SO3.
Increasing the temperature: This reaction is exothermic, meaning it releases heat. Increasing the temperature would favor the reverse reaction, consuming SO3 and forming SO2.
Decreasing the pressure: The reaction involves a decrease in moles (from 2 to 1), so decreasing pressure would favor the forward reaction, producing more SO3 to increase pressure.
Therefore, adding more SO3 is the only option that would prevent the reaction from shifting towards forming more of the product.