In a reversible reaction, a catalyst (a) increases the rate of forward reaction only (b) increases the rate of forward reaction to a greater extent that of the backward reaction (c) increases the rate of forward reaction and decreases that of the backward reaction (d) increases the rate of forward and backward reaction by the same factor.

Now, when students approach problems involving equilibrium, it's crucial to consider not just the state at which the reaction sits but how it arrived there and how changes will affect it. For instance, increasing the temperature may favor one side of the equilibrium—and understanding this can help students predict and control chemical reactions effectively.

It is vital to remember, though, that activation energy is just that—an initial hurdle. Lowering it does not alter the final balance of reactants and products, nor does it influence how much energy is released or absorbed in the reaction overall. It simply makes reaching equilibrium faster.

A catalyst is often used to increase the rate of a reaction, making it a central point of interest, especially in industrial processes where time is money. By lowering the activation energy, as discussed earlier, a catalyst permits more particles to react over the same period, thus speeding up both the forward and backward reactions in a reversible process. Nonetheless, despite hastening the journey to equilibrium, the catalyst does not change the position of the equilibrium or the proportions of reactants and products at equilibrium, ensuring that the chemistry remains unchanged even as the tempo changes.

Understanding these principles allows students to manipulate reaction conditions for desired outcomes in chemical processes, be it in a laboratory or commercial setting.