Final answer:
The true statement about chemical reaction kinetics is that the activation energy determines the reaction speed. It is the minimum energy required to initiate a reaction. Higher activation energy leads to a slower reaction rate.
Step-by-step explanation:
The true statement about the kinetics of a chemical reaction is c) Activation energy determines reaction speed. The relationship between reaction rate and concentration is represented as a rate law, and the rate indeed depends on the concentration of the reactants, disputing choice a). Catalysts, contrary to b), actually speed up reactions by providing an alternate pathway with a lower activation energy. Equilibrium is a state that is reached over time, not instantaneously, making d) incorrect.
Essentially, activation energy is the minimum energy required for a reaction to occur. When reactants have sufficient energy to overcome the activation energy barrier, they form the transition state, then products. Higher activation energy means fewer molecules have enough energy to react, leading to a slower reaction rate.
Moreover, reactant collisions must involve proper orientation and enough kinetic energy to surpass the activation energy. Therefore, increased temperature - increasing the reactants' kinetic energy - commonly leads to a higher reaction rate.