Final answer:
Controlling the side equilibrium favors in an acid-catalyzed hydration reaction involves applying Le Chatelier's principle. Adjusting concentrations of reactants and products can shift equilibrium towards either side. Removing a product or adding excess reactants drives the forward reaction, while adding excess products or removing reactants favors the reverse reaction.
Step-by-step explanation:
To control which side equilibrium will favor in an acid-catalyzed hydration reaction, one must consider Le Chatelier's principle. This principle states that if a system at equilibrium is subjected to change in concentration, temperature, volume, or pressure, the system will adjust to counteract the change and re-establish equilibrium. In the context of a reaction, manipulating concentrations of reactants or products can shift the equilibrium. For instance:
- Adding more of a reactant typically causes the equilibrium to shift to the product side, favoring the forward reaction.
- Conversely, removing a product as it is formed will also drive the reaction forward.
- If you add more of a product, the equilibrium shifts to the reactant side, which means the reverse reaction is favored.
- Removing a reactant will similarly cause the reaction to favor the reverse process to create more of the removed reactant.
These changes in concentration will affect the reaction's direction and the overall concentrations of reactants and products at equilibrium. For example, if water is removed from the products in an acid-catalyzed dehydration, the reaction can be driven to completion, making it irreversible. Conversely, adding excess water can provoke the reverse reaction.