Enthalpy and entropy are two fundamental concepts in thermodynamics that describe energy changes and disorder in chemical reactions.

Enthalpy is a thermodynamic property that quantifies the heat energy released or absorbed during a chemical reaction. It is denoted by the symbol $\Delta H$ and is typically measured in kilojoules per mole ($\text{kJ/mol}$). If $\Delta H$ is negative, the reaction releases heat energy, classifying it as exothermic. Conversely, a positive $\Delta H$ indicates that the reaction absorbs heat energy, making it endothermic.

Entropy, represented by the symbol $\Delta S$, measures the degree of disorder or randomness within a system. This property is usually expressed in joules per mole per Kelvin ($\text{J/(mol·K)}$). A positive value for $\Delta S$ signifies an increase in disorder, while a negative value suggests that the system becomes more ordered.

The interplay between enthalpy and entropy is articulated through the Gibbs free energy equation. This equation states that the change in free energy ($\Delta G$) is given by:

where $T$ is the absolute temperature in Kelvin. A negative value for $\Delta G$ indicates that the reaction is spontaneous, meaning it can proceed without the need for additional energy input.

In summary, enthalpy quantifies the heat energy involved in chemical reactions, while entropy measures the level of disorder in a system. Their relationship, as defined by the Gibbs free energy equation, is crucial for determining the spontaneity of reactions.