Differences between Endothermic and Exothermic
Contents
Endothermic vs. Exothermic Processes
In thermochemistry, chemical reactions are classified as either endothermic or exothermic based on the flow of heat. An exothermic process releases energy into the surroundings, typically as heat, causing the temperature of the surroundings to rise.[1][2] Conversely, an endothermic process absorbs heat from its surroundings, leading to a decrease in the temperature of the surroundings.[3][4]
This exchange of energy is a result of the breaking and forming of chemical bonds during a reaction.[5] Energy is required to break the bonds within reactant molecules, an endothermic process. Energy is released when new bonds are formed in the product molecules, which is an exothermic process. The net energy change of the reaction determines whether it is classified as endothermic or exothermic.[5] If more energy is released in bond formation than is used for bond breaking, the overall reaction is exothermic.[5] If more energy is absorbed to break bonds than is released during formation, the reaction is endothermic.[5]
The change in enthalpy (ΔH) is a measure used to quantify this heat exchange between the reaction (the system) and its surroundings. Exothermic reactions have a negative enthalpy change (ΔH < 0), indicating a net release of energy. Endothermic reactions have a positive enthalpy change (ΔH > 0), signifying a net absorption of energy.
[4]== Comparison Table ==
| Category | Endothermic | Exothermic |
|---|---|---|
| Energy Flow | Absorbs heat from surroundings | [2] Releases heat into surroundings |
| Temperature of Surroundings | Decreases; feels cold | [1] Increases; feels warm |
| Enthalpy Change (ΔH) | Positive (ΔH > 0) | Negative (ΔH < 0) |
| Bond Energy | Energy needed to break bonds is greater than energy released forming bonds | Energy released forming bonds is greater than energy needed to break bonds |
| [5]Reactant vs. Product Energy | Products have higher energy than reactants | Products have lower energy than reactants |
| Common Physical Examples | Melting ice, boiling water, evaporation | Freezing water, condensation of rain |
| [1]Common Chemical Examples | Photosynthesis, cooking an egg, dissolving ammonium nitrate in water | Combustion (burning), respiration, neutralization of acid and base |
Spontaneity of Reactions
The spontaneity of a reaction is not solely determined by its enthalpy change. While[3] many exothermic reactions occur spontaneously, endothermic reactions can also be spontaneous. Spontaneity is determined by the change in Gibbs free energy, which accounts for both enthalpy and the change in entropy (a measure of disorder). An endothermic[3] reaction can proceed spontaneously if it is accompanied by a significant increase in entropy.
Examples[3] in Daily Life
Exothermic and endothermic reactions are common in everyday life. Combustion reactions, such as burning wood or fuel in a vehicle, are exothermic processes that release significant energy. Respiration in living organisms is another example of a complex exothermic process where glucose is broken down to release energy. Neutralization reactions between acids and bases are also typically exothermic.
Common endothermic processes include physical changes like melting ice cubes and evaporating water, where energy is absorbed from the environment to cause the phase change. Photosynthesis is a vital endothermic chemical reaction in which plants use energy from sunlight to convert carbon dioxide and water into glucose and oxygen. The application of a cold pack to an injury often utilizes an endothermic reaction, such as dissolving ammonium nitrate in water, which absorbs heat from the surroundings and creates a cooling effect.
