Final answer:
Lowering the activation energy of a spontaneous reaction would typically lead to an increase in the reaction rate. This is because a lower activation energy allows more reactant molecules to have sufficient energy to successfully collide and react, thereby speeding up the reaction.
Step-by-step explanation:
If the activation energy of a spontaneous reaction were lowered, the rate of the reaction would typically increase. Activation energy is defined as the minimum amount of energy that reactant molecules must possess for a reaction to occur after a successful collision. Lowering the activation energy makes it easier for reactant molecules to overcome the energy barrier, leading to an increase in the number of successful collisions per unit of time. Subsequently, this increases the rate at which reactants are converted into products.
For example, catalysts are substances that lower the activation energy without being consumed in the reaction, thus, they can significantly increase the reaction rate. A catalyst provides an alternate pathway for the reaction with a lower activation energy. An example can be seen in biological enzymes, which are natural catalysts that increase the rates of chemical reactions in living organisms by lowering their activation energies.