An exothermic reaction releases heat, while entropy measures a system’s disorder. Although often linked, these concepts describe fundamentally different aspects of a chemical process: exothermicity focuses on the energy change within the system and its surroundings, while entropy describes the distribution of matter and energy. Determining the outcome of a chemical change requires considering both forces, as they do not always move in the same direction.
The Energy Factor: What Exothermic Means
An exothermic reaction releases energy, typically as heat, into its surroundings. This release is quantified by the change in enthalpy (\(\Delta H\)), which is negative for all exothermic processes. The chemical energy stored in the product bonds is lower than the energy stored in the initial reactant bonds, and this difference is the heat expelled into the environment.
For example, the combustion of natural gas (methane) is strongly exothermic. This process is used to heat homes because the products, carbon dioxide and water, hold significantly less energy than the reactants. The negative enthalpy change indicates that the system has lost internal energy, which is a driving force that favors a reaction, as chemical systems tend toward lower energy states.
The Disorder Factor: Defining Entropy
Entropy, symbolized as \(\Delta S\), measures the dispersal of energy and matter within a system. It reflects the number of different ways the energy and molecules can be arranged; a system with a high number of possible arrangements (microstates) has high entropy.
Entropy increases when a system moves from a more ordered state to a less ordered one. For instance, converting a solid to a liquid or a liquid to a gas results in a large increase in entropy because the molecules gain freedom of movement. Similarly, a chemical reaction that increases the number of gas molecules from reactants to products will generally have a positive change in entropy.
The second law of thermodynamics states that the total entropy of the universe must increase for any spontaneous process to occur. An increase in a system’s entropy is a fundamental driving force that favors a chemical change.
How Exothermicity and Entropy Determine Reaction Outcomes
The spontaneity of a chemical reaction is determined by the combined influence of enthalpy and entropy. These two factors are mathematically linked through the concept of Gibbs Free Energy (\(\Delta G\)), which represents the energy available in a system to do useful work. A reaction is considered spontaneous if the change in Gibbs Free Energy is negative.
The relationship is expressed as \(\Delta G = \Delta H – T\Delta S\), where \(T\) is the absolute temperature. The equation shows that a negative change in enthalpy (\(\Delta H\), an exothermic reaction) favors spontaneity by contributing to a negative \(\Delta G\). Likewise, a positive change in entropy (\(\Delta S\), increasing disorder) also contributes to a negative \(\Delta G\).
Because these two factors can reinforce or oppose each other, an exothermic reaction releasing heat transfers energy to the surroundings, increasing the entropy of the environment. Therefore, even if the system’s own entropy decreases, the overall entropy of the universe can still increase, allowing the reaction to be spontaneous.
Scenarios Where Heat Release and Disorder Interact
The answer to whether an exothermic reaction increases entropy depends on which part of the system is being measured. The entropy of the chemical system itself may increase, decrease, or stay roughly the same.
A combustion reaction, like burning wood, is highly exothermic and increases the system’s entropy because it converts a solid and a gas into a larger number of gas molecules. This combination of negative \(\Delta H\) and positive \(\Delta S\) makes the reaction spontaneous at all temperatures.
Conversely, some exothermic processes decrease the system’s entropy. For example, the freezing of water is exothermic, releasing heat as liquid water turns into solid ice. This process decreases entropy as mobile water molecules arrange into an ordered crystal structure. In this case, the negative \(\Delta H\) is large enough to overcome the negative \(\Delta S\), but only at low temperatures (below \(0^\circ \text{C}\)). While exothermicity makes a reaction more likely to happen, it does not guarantee an increase in the system’s entropy; it only guarantees an increase in the entropy of the surroundings.