As the reaction \(\mathrm{A} \rightarrow \mathrm{B}\) proceeds, the container in which it is run feels hot to the touch. Measurements show a net energy change of \(250 \mathrm{~kJ} / \mathrm{mol}\). (a) What is \(\Delta E_{\text {forward rxn }}\) ? (b) What is \(\Delta E_{\text {reverse rxn }}\) ?

Understanding the heat of chemical reactions is fundamental in the study of chemistry. When a chemical reaction occurs, energy can be either absorbed or released into the surrounding environment. An exothermic reaction is a type of chemical reaction that releases energy by light or heat. This is the type of reaction that might make a container feel hot to the touch, signifying that heat is being produced and dispersed into the surroundings.

Common examples of exothermic reactions include combustion, like a fire burning wood, and many types of chemical explosives when they detonate. In terms of enthalpy, an exothermic reaction has a negative enthalpy change because the energy of the products is less than the energy of the reactants, indicating that energy has been released into the surroundings. The notation for this energy change is a negative sign before the value, such as \( \-250~\mathrm{kJ/mol} \) in the provided exercise.

Enthalpy, represented by the symbol \( H \) and sometimes referred to as heat content, is a measure of the total energy of a thermodynamic system. When a chemical reaction occurs, the enthalpy change, denoted as \( \Delta H \) for a reaction, indicates whether the reaction absorbs heat from the environment (endothermic) or releases heat into the environment (exothermic).

The value of \( \Delta H \) is calculated by subtracting the enthalpy of the reactants from the enthalpy of the products. In the context of the exercise provided, the enthalpy change during the exothermic reaction, where product \( \mathrm{B} \) is formed from reactant \( \mathrm{A} \) is \( \Delta E_{\text{forward rxn}} = -250~\mathrm{kJ/mol} \) indicating the release of energy. This is often illustrated in a diagram showing a downward slope from reactants to products, representing the energy decrease.

The concept of reverse reaction energy is directly tied to the principle of reversibility in chemical reactions. In a simple reaction represented as \( \mathrm{A} \rightarrow \mathrm{B} \) with an exothermic forward reaction, the reverse reaction \( \mathrm{B} \rightarrow \mathrm{A} \) would necessarily be endothermic.

During the reverse process, energy must be absorbed from the surroundings to convert the products back into reactants. Consequently, the sign of the energy change for the reverse reaction is positive. This is because energy is gained by the system from the surroundings. In the example exercise, the reverse reaction energy \( \Delta E_{\text{reverse rxn}} \) is simply the opposite of the forward reaction energy, resulting in \( +250~\mathrm{kJ/mol} \), which demonstrates the amount of energy absorbed to revert product \( \mathrm{B} \) back into reactant \( \mathrm{A} \). It's crucial to understand that the magnitude of energy remains the same; only its direction (release or absorption) changes with the direction of the reaction.