Question

The enthalpy change for a reaction, ?H, is negative. What does this indicate about the chemical potential energy of the system before and after the reaction?

Short answer

A negative enthalpy change, \(\Delta H\), indicates an exothermic reaction where the chemical potential energy of the reactants (before the reaction) is higher than that of the products (after the reaction). As the reaction proceeds, energy is released in the form of heat because the system's chemical potential energy decreases.

Step-by-step solution

Understanding Enthalpy Change

Enthalpy change (?H) represents the heat transferred during a reaction under constant pressure. A positive ?H means heat is absorbed by the system (endothermic reaction), while a negative ?H means heat is released by the system (exothermic reaction).

Chemical Potential Energy Before and After the Reaction

In a reaction, the chemical potential energy is related to the energy stored in the chemical bonds of the reactants and products. To understand how the chemical potential energy of the system changes during a reaction, we can think of the energy of the bonds in the reactants and the energy of the bonds in the products.

Interpreting Negative Enthalpy Change

Since the enthalpy change (?H) in this exercise is negative, this indicates an exothermic reaction. In an exothermic reaction, the chemical potential energy of the reactants (before the reaction) is higher than the chemical potential energy of the products (after the reaction). Therefore, energy is released during the reaction in the form of heat.

Conclusion

A negative enthalpy change (?H) in a reaction indicates that the chemical potential energy of the system decreases during the reaction. The reactants, which have higher chemical potential energy prior to the reaction, release energy as they form products with lower chemical potential energy. This energy is released in the form of heat because the reaction is exothermic.

Key concepts

Exothermic Reaction

An exothermic reaction is a fascinating process where energy is released, usually in the form of heat, by the reaction. This type of reaction is identified by a negative enthalpy change (\( \Delta H \)), indicating that the system releases energy to its surroundings.

Exothermic reactions are commonly observed in everyday life, such as during combustion processes, like burning wood in a fireplace. When wood burns, it releases stored energy in the form of heat and light, warming up our surroundings. Similarly, many chemical reactions, such as the rusting of iron, also fall under this category.

These reactions are characterized by a decrease in chemical potential energy from reactants to products, which results in the release of excess energy. The system's temperature may rise, depending on the surroundings, making exothermic reactions significant for both natural processes and industrial applications.

Chemical Potential Energy

Chemical potential energy is the stored energy within the chemical bonds of a substance. It plays a crucial role in understanding how energy is utilized and transformed in reactions.

Every molecule consists of atoms held together by bonds, and these bonds contain varying amounts of energy depending on their type and strength. Before a reaction occurs, the reactants possess a certain amount of chemical potential energy derived from their specific bonding arrangements. This energy is what dictates how a particular reaction will transform:

• Reactions involving breaking strong bonds will require more energy.
• Forming strong bonds will release more energy.

In an exothermic reaction, the energy stored within the reactants' chemical bonds is greater than the energy in the products' bonds. As the reaction progresses, the excess energy is released, resulting in products with a lower chemical potential energy. This decrease signifies the energy transformation from potential to kinetic, often observed as heat energy released to the environment.

Constant Pressure Reaction

A constant pressure reaction is a chemical reaction that occurs under a condition where the pressure remains unchanged throughout the process.

Most reactions we encounter in everyday situations, such as burning or dissolving substances, happen under constant pressure—generally at atmospheric pressure. This setting is beneficial as it simplifies the study of thermochemical processes by focusing on enthalpy changes without worrying about pressure-induced volume changes.

During such reactions, we focus on the enthalpy change (\( \Delta H \)) to determine whether the process is endothermic or exothermic. Since the exercise highlights a negative \( \Delta H \), it tells us that under these constant pressure conditions, the system is losing energy as heat to its environment. Understanding that many reactions naturally occur at constant pressure helps better relate theoretical concepts to real-world applications, such as optimizing energy outputs in combustion engines and other industrial exothermic reactions.