Suppose that the reaction A->products is exothermic and has an activation barrier of 75 kJ>mol. Sketch an energy diagram showing the energy of the reaction as a function of the progress of the reaction. Draw a second energy curve showing the effect of a catalyst.

Understanding activation energy helps explain why certain reactions occur and why some do not. It refers to the minimum amount of energy that reacting particles must possess for a chemical reaction to take place. This energy is necessary to break the bonds within reactants, leading to a chemical transformation. In other words, activation energy can be imagined as a 'hill' that reactants need to climb so they can roll down the other side as products.

In an energy diagram, this is represented as a peak that must be overcome from the energy level of the reactants to the highest point of the curve before descending to the products’ energy level. A higher peak means a higher activation energy, making it harder for the reaction to proceed. Reactant A, with a noted activation barrier of 75 kJ/mol, would thus require a significant input of energy to reach the transition state, a concept termed as the 'activated complex,' before the reaction can go downhill towards product formation.

Exothermic reactions are chemical processes that release energy in the form of heat or light. In the context of an energy diagram, these reactions can be identified when the energy level of the products is lower than that of the reactants. This type of reaction is favorable from an energy standpoint because it results in a net release of energy to the surroundings.

Since reactant A to products is described as an exothermic process, the energy curve starts high with reactant A, climbs up to the activation energy peak, and then falls to a lower energy level for the products. This downward slope signifies that more energy is released then was initially required to reach the activation peak. The difference in energy (horizontal gap) between the reactant and product is effectively the energy given off during the reaction, which can be observed as heat or sometimes light.

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process. It works by providing an alternative pathway for the reaction with a lower activation energy. On an energy diagram, this is visualized by a process that begins at the same energy level but has a notably lower ‘hill’ or peak to surpass.

The effect of a catalyst is pivotal as it makes reactions more accessible and faster by reducing the activation energy barrier. For example, if a catalyst is introduced into the reaction of A to products, the energy curve would depict a lower peak than the original 75 kJ/mol needed in the uncatalyzed reaction. It’s important to remember that while catalysts accelerate reactions, they do not alter the initial or final energy levels of the reactants and products; they merely make it easier for the reaction to proceed by lowering the energy hill.