Exothermic reaction
Chemical reaction releasing energy as light or heat
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Examples
Examples are numerous: combustion, the thermite 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:
- CH4 + 2O2 → CO2 + 2H2O ΔH⚬ = −890 kJ/mol
These sample reactions are strongly exothermic.
Uncontrolled exothermic reactions, those leading to fires and explosions, are wasteful because it is difficult to capture the released energy. Nature effects combustion reactions under highly controlled conditions, avoiding fires and explosions, in aerobic respiration so as to capture the released energy, e.g. for the formation of ATP.
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Measurement
The enthalpy of a chemical system is essentially its energy. The enthalpy change ΔH for a reaction is equal to the heat q 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 using calorimetry, e.g. with a bomb calorimeter. One common laboratory instrument is the reaction calorimeter, where the heat flow from or into the reaction vessel is monitored. The heat release and corresponding energy change, ΔH, of a combustion reaction can be measured particularly accurately.
The measured heat energy released in an exothermic reaction is converted to ΔH⚬ in Joule per mole (formerly cal/mol). The standard enthalpy change ΔH⚬ is essentially the enthalpy change when the stoichiometric 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 to bond energies to a good approximation by:
- ΔH⚬ = (total bond energy of reactants) − (total bond energy of products)
In an exothermic reaction, by definition, the enthalpy change has a negative value:
- ΔH = Hproducts − Hreactants < 0
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:
- 2H2 (g) + O2 (g) → 2H2O (g)
- ΔH⚬ = −483.6 kJ/mol