Each of the statements given below is false. Explain why. a. The activation energy of a reaction depends on the overall energy change \((\Delta E)\) for the reaction. b. The rate law for a reaction can be deduced from examination of the overall balanced equation for the reaction. c. Most reactions occur by one-step mechanisms.

Statement a: The activation energy of a reaction and the overall energy change \((\Delta E)\) are not directly related. Activation energy is related to the pathway of a chemical reaction, whereas the overall energy change \((\Delta E)\) is related to the initial and final states of the system. Statement b: The rate law cannot be determined directly from the balanced chemical equation, as it depends on the elementary steps and the mechanism of the reaction. The rate law is experimentally determined by measuring the reaction rate with different initial concentrations of reactants. Statement c: Most reactions actually proceed through a series of elementary steps or multiple-step mechanisms, as one-step mechanisms are rare and require simultaneous collisions among all reactants with correct orientation and energy. In reality, many reactions occur through a sequence of simpler steps leading to the formation of products.