Final answer:
A catalyst lowers the activation energy of a reaction by providing a new pathway and stabilizing the transition state, making it easier for reactants to turn into products, which increases the reaction rate.
Step-by-step explanation:
When a catalyst lowers the activation energy of a reaction, it provides a new reaction pathway with a lower activation energy. This does not mean that it provides energy to the reactants (A) or increases the activation energy (B). Instead, in molecular terms, it means that the catalyst stabilizes the transition state (C), reducing the amount of energy required for the reaction to proceed. It also doesn't mean the potential energy of the reactants is increased (D); the energies of the reactants and products remain unchanged by the presence of the catalyst.
The mechanism by which this happens often involves the catalyst offering a surface or an environment that brings reactants closer, aligns them suitably for a reaction, or provides a site where reactants can be temporarily bound to alter bond angles and electron distribution, all contributing to lowering the activation energy.
An example is the action of enzymes in biochemical reactions that bind substrates and bring them into close proximity at an active site, facilitating the reaction without being consumed in the process. The reaction rate increases because more reactant molecules are able to undergo effective collisions and form products as a result of the lowered activation energy.