Inthermochemistry, anexothermic reaction is a "reaction for which the overallstandard enthalpy change ΔH⚬ is negative."[1][2] Exothermic reactions usually releaseheat. The term is often confused withexergonic reaction, which IUPAC defines as "... a reaction for which the overall standard Gibbs energy change ΔG⚬ is negative."[2] A strongly exothermic reaction will usually also be exergonic because ΔH⚬ makes a major contribution toΔG⚬. Most of the spectacular chemical reactions that are demonstrated in classrooms are exothermic and exergonic. The opposite is anendothermic reaction, which usually takes up heat and is driven by anentropy increase in the system.
Examples are numerous:combustion, thethermite reaction, combining strong acids and bases,polymerizations. As an example in everyday life,hand warmers make use of the oxidation of iron to achieve an exothermic reaction:
A particularly important class of exothermic reactions is combustion of a hydrocarbon fuel, e.g. the burning of natural gas:
These sample reactions are strongly exothermic.
Uncontrolled exothermic reactions, those leading tofires andexplosions, are wasteful because it is difficult to capture the released energy. Nature effects combustion reactions under highly controlled conditions, avoiding fires and explosions, inaerobic respiration so as to capture the released energy, e.g. for the formation ofATP.
Theenthalpy of a chemical system is essentially its energy. The enthalpy change ΔH for a reaction is equal to the heatq transferred out of (or into) a closed system at constant pressure without in- or output of electrical energy. Heat production or absorption in a chemical reaction is measured usingcalorimetry, e.g. with abomb calorimeter. One common laboratory instrument is thereaction calorimeter, where the heat flow from or into the reaction vessel is monitored. The heat release and corresponding energy change, ΔH, of acombustion reaction can be measured particularly accurately.
The measured heat energy released in an exothermic reaction is converted to ΔH⚬ inJoule per mole (formerlycal/mol). Thestandard enthalpy change ΔH⚬ is essentially the enthalpy change when thestoichiometric coefficients in the reaction are considered as the amounts of reactants and products (in mole); usually, the initial and final temperature is assumed to be 25 °C. For gas-phase reactions, ΔH⚬ values are related tobond energies to a good approximation by:
In an exothermic reaction, by definition, the enthalpy change has a negative value:
where a larger value (the higher energy of the reactants) is subtracted from a smaller value (the lower energy of the products). For example, when hydrogen burns:
