Suppose that the reaction \(\mathrm{A} \longrightarrow\) products is exothermic and has an activation barrier of \(75 \mathrm{~kJ} / \mathrm{mol} .\) 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.

Exothermic reactions are chemical processes that release energy into their surroundings, typically in the form of heat. These reactions result in a higher temperature of the surrounding environment. The classic indicators of an exothermic reaction include the warmth felt from an exothermic handwarmer or the heat produced when wood burns.

On an energy diagram, an exothermic reaction is represented by showing the products at a lower energy level than the reactants. This conveys that energy has been released to their surroundings as the reaction progresses. The energy difference between the reactants and products is called the heat of reaction or enthalpy change (\text{\(\tlDelta H\)}). For example, if a reaction starts with reactants at a certain energy level and ends with products at a lower level, the difference in these energy levels is the heat released during the reaction. In educational illustrations, the reactants are often placed at a higher point on the energy axis, sloping downward to the products, demonstrating the decrease in internal energy.

Activation energy (\text{\(E_a\)}) is the minimum energy needed to initiate a chemical reaction. It can be envisioned as the 'energy barrier' that reactants must overcome for a reaction to occur. The concept of activation energy is crucial in understanding why some reactions occur spontaneously while others require input of energy, such as heat or light, to proceed.

An energy diagram visualizes this barrier as a peak that the reactants must climb in order to transform into products. Once the reactants have absorbed enough energy to reach the top of the peak, they undergo a transformation, passing through a high-energy, unstable transition state before releasing energy and becoming the lower-energy products in an exothermic reaction. The activation energy is depicted on the diagram as the height difference between the reactants' initial energy level and the peak of the curve.

Catalysts play a pivotal role in facilitating chemical reactions without being consumed in the process. In chemistry, a catalyst provides an alternative pathway for a reaction with a lower activation energy, thereby increasing the reaction rate. Catalysts are incredibly useful in industrial processes, biological systems, and laboratory reactions because they can dramatically reduce the energy requirements and expedite reactions that would otherwise be impractically slow or require high-energy conditions.

On an energy diagram, the effect of a catalyst is shown by comparing two curves: one without the catalyst and one with it. The latter curve peaks at a lower point, effectively lowering the 'hill' that the reactants must climb. While the start and end points remain the same—indicating that catalysts do not change the overall energy change of the reaction—the peak representing the transition state is lowered, illustrating the reduced activation energy required to reach the transition state thanks to the catalyst.

The transition state, also known as the activated complex, represents a fleeting configuration of atoms at the point of highest energy during a chemical reaction. It is a state of maximum instability where old bonds are breaking, and new bonds are forming. The transition state cannot be isolated because of its short-lived existence, and it occupies a critical point on the energy diagram, marking the peak of the curve that represents the activation energy barrier.

In exothermic reactions, the energy diagram highlights the transition state as a peak which reactants must reach before descending towards the products. It signifies a decisive point in the reaction path where the potential energy is at its maximum and the fate of the reactants—to revert to the original state or to proceed to form products—is determined. This is a crucial concept in chemical kinetics, as the nature and energy of the transition state dictate the rate at which the reaction will proceed.