As an expert in the field of thermodynamics, I can provide you with a comprehensive understanding of the relationship between the sign of the heat of reaction (denoted as \( \Delta H \), where "H" stands for enthalpy) and whether a reaction is exothermic or endothermic.

**Step 1: Understanding Enthalpy and the Sign Convention**
Enthalpy is a thermodynamic property that represents the total heat content of a system. It is defined as the sum of the internal energy of the system and the product of its pressure and volume. In chemical reactions, the change in enthalpy (\( \Delta H \)) is the difference between the enthalpy of the products and the enthalpy of the reactants.

The sign convention for \( \Delta H \) is such that if the system releases heat to the surroundings (which is the case for an exothermic reaction), \( \Delta H \) is negative. Conversely, if the system absorbs heat from the surroundings (an endothermic reaction), \( \Delta H \) is positive.

**Step 2: The Nature of Exothermic Reactions**
Exothermic reactions are characterized by the release of energy in the form of heat. This release occurs because the energy of the products is lower than the energy of the reactants. When the products have less energy than the reactants, the excess energy is given off as heat, resulting in a negative \( \Delta H \).

**Step 3: The Sign of \( \Delta H \) for Exothermic Reactions**
Given the sign convention and the nature of exothermic reactions, it is clear that \( \Delta H \) is negative for exothermic reactions. This is because the enthalpy of the products is lower than that of the reactants, and the system loses energy to the surroundings in the form of heat.

Step 4: Examples of Exothermic Reactions
Many common processes are exothermic, such as:
- Combustion reactions, where a substance reacts with oxygen to release heat.
- Neutralization reactions, where an acid and a base react to form water and a salt, releasing heat in the process.
- Crystallization, where a substance transitions from a less ordered state (such as a liquid or gas) to a more ordered state (such as a solid), releasing heat.

**Step 5: The Importance of Understanding \( \Delta H \)**
Understanding the sign of \( \Delta H \) is crucial for predicting the spontaneity of a reaction, determining the conditions under which a reaction will occur, and designing industrial processes that can harness the energy released during exothermic reactions.

In conclusion, for an exothermic reaction, \( \Delta H \) is negative because the system releases heat to the surroundings, and the enthalpy of the products is lower than that of the reactants.

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A system that releases heat to the surroundings, an exothermic reaction, has a negative --H by convention, because the enthalpy of the products is lower than the enthalpy of the reactants of the system.Mar 8, 2017read more >>