Exothermic reaction
A reaction with negative standard enthalpy change, releasing heat.
An exothermic reaction is a chemical reaction for which the overall standard enthalpy change ΔH⚬ is negative. Such reactions usually release stored energy as heat and are often confused with exergonic reactions, which are defined by a negative standard Gibbs energy change ΔG⚬. A strongly exothermic reaction will usually also be exergonic because ΔH⚬ makes a major contribution to ΔG⚬.
- field
- Thermochemistry
- known_for
- Reactions with negative standard enthalpy change, releasing heat
- opposite
- Endothermic reaction
Lore & Background
Exothermic reactions are defined in thermochemistry as reactions for which the overall standard enthalpy change ΔH⚬ is negative. They usually release stored energy as heat. The term is often confused with exergonic reaction, which IUPAC defines as a reaction for which the overall standard Gibbs energy change ΔG⚬ is negative. A strongly exothermic reaction will usually also be exergonic because ΔH⚬ makes a major contribution to ΔG⚬. Most spectacular classroom chemical demonstrations are both exothermic and exergonic. The opposite is an endothermic reaction, which usually takes up heat and is driven by an entropy increase in the system. Examples include combustion, the thermite reaction, combining strong acids and bases, and polymerizations. Uncontrolled exothermic reactions leading to fires and explosions are wasteful because it is difficult to capture the released energy. Nature effects combustion reactions under highly controlled conditions in aerobic respiration to capture energy for ATP formation.
Reader's Guide
Exothermic reactions are fundamental in thermochemistry, defined by a negative standard enthalpy change (ΔH⚬). Their significance lies in the release of stored energy as heat, which powers many natural and industrial processes. The measurement of heat release uses calorimetry, such as a bomb calorimeter or reaction calorimeter, with the enthalpy change expressed in joules per mole. The standard enthalpy change ΔH⚬ is essentially the enthalpy change when stoichiometric coefficients are considered as amounts in moles, typically at 25 °C. For gas-phase reactions, ΔH⚬ relates to bond energies: ΔH⚬ = (total bond energy of reactants) − (total bond energy of products). The negative value indicates that the enthalpy of products is lower than that of reactants. This concept is central to understanding combustion, respiration, and many chemical reactions. The distinction from exergonic reactions (based on Gibbs energy) is important, though strongly exothermic reactions are usually also exergonic. The legacy includes applications from hand warmers to controlled biological energy capture, highlighting both the utility and hazards of uncontrolled exothermic processes.
Did You Know?
- An exothermic reaction is defined by a negative standard enthalpy change ΔH⚬.
- The term is often confused with exergonic reaction, which has a negative standard Gibbs energy change ΔG⚬.
- A strongly exothermic reaction will usually also be exergonic because ΔH⚬ makes a major contribution to ΔG⚬.
- Uncontrolled exothermic reactions leading to fires and explosions are wasteful because it is difficult to capture the released energy.
Frequently Asked Questions
Who is Exothermic reaction?
In the Physical Chemistry And Thermodynamics 1-20 canon, Exothermic reaction is the thermochemistry entry defined by a negative standard enthalpy change (ΔH⚬ < 0). It represents any chemical process in which the products sit at a lower enthalpy than the reactants, so the excess energy is expelled to the surroundings, most often as heat.
What are Exothermic reaction's powers/role?
Its signature move is dumping stored chemical energy into the environment as thermal energy, which is why a test tube or reactor gets noticeably warm during the process. Fans often mix it up with the Exergonic reaction, but the two are governed by different state functions—enthalpy versus Gibbs free energy.
How does Exothermic reaction's story end?
A strongly exothermic process almost always wraps up as an exergonic one too, because the large negative ΔH⚬ dominates the ΔG⚬ = ΔH⚬ − TΔS⚬ relationship and pushes the reaction forward spontaneously. In weaker cases, a positive entropy term can tip the balance, so exothermic does not strictly guarantee exergonic.
Why is Exothermic reaction important?
It sits at the core of thermochemistry and explains why combustion, acid-base neutralization, and countless industrial syntheses release palpable heat. Grasping its negative ΔH⚬ is the foundational step for predicting energy flow, designing calorimetry experiments, and sizing safety margins in real plants.
What is Exothermic reaction's rival?
Its direct counterpart in the canon is the Endothermic reaction, which carries a positive standard enthalpy change and absorbs heat from the surroundings instead of releasing it. The two are mirror images in the enthalpy ledger, and comparing them is the standard way textbooks introduce the first law of thermodynamics.
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