Final answer:
The correct option that will cause a reaction at equilibrium to go in the reverse direction, forming more reactants, is b) Removing a reactant from the reaction container. The correct answer is option b.
Step-by-step explanation:
To understand what causes a reaction at equilibrium to proceed in the reverse direction, we must consider Le Chatelier's principle, which describes how a system at equilibrium responds to changes in concentration, pressure, or temperature. Specifically, the reversal of a reaction so that products revert to reactants can occur under certain conditions.
Effect of Changes in Concentration
When we add more of a reactant (option a) to a reaction at equilibrium, Le Chatelier's principle tells us that the system will shift to counter this by increasing the rate of the forward reaction to consume the added reactant, resulting in more products. On the other hand, if a reactant is removed (option b), the reaction will shift to replace the reduced concentration of the reactant, effectively favoring the reverse reaction. Similarly, when a product is removed (option c), the equilibrium is also disturbed, but in this case, to counter the decrease in product concentration, the reaction will favor the formation of more products from the reactants, thus increasing the forward reaction.
Therefore, the correct option that will cause the equilibrium reaction to go in the reverse direction is b) Removing a reactant from the reaction container. This is because the system seeks to restore equilibrium by shifting the reaction towards the left, to make more of the removed reactant.
Real-World Example: The Haber Process
In the industrial Haber-Bosch process, the removal of NH3 (ammonia) shifts the reaction to produce more NH3 from nitrogen and hydrogen gases. If N₂ or H₂ is removed, the reverse reaction is favored to replace the depleted gas, which shows the principle in action.