Editing Exothermic reaction

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{{short description|Chemical reaction that releases energy as light or heat}}
[[Image:ThermiteReaction.jpg|thumb|The [[thermite]] reaction is famously exothermic. The reduction of iron(III) oxide by [[aluminium]] releases sufficient heat to yield molten iron.]]
In [[thermochemistry]], an '''exothermic reaction''' is a "reaction for which the overall [[Standard enthalpy of reaction|standard enthalpy change]] Δ''H''⚬ is negative."<ref>{{cite book |chapter=Exothermic reaction |chapter-url=https://goldbook.iupac.org/terms/view/E02269 |publisher=IUPAC|doi=10.1351/goldbook.E02269 |title=The IUPAC Compendium of Chemical Terminology |year=2014 }}</ref><ref name="IUPAC Recommendations 1996">{{cite journal |doi=10.1351/pac199668010149|title=A glossary of terms used in chemical kinetics, including reaction dynamics (IUPAC Recommendations 1996)|year=1996|last1=Laidler|first1=K. J.|s2cid=98267946|journal=Pure and Applied Chemistry|volume=68|pages=149–192|doi-access=free}}</ref> Exothermic reactions usually release [[heat]]. The term is often confused with [[exergonic reaction]], which IUPAC defines as "... a reaction for which the overall standard Gibbs energy change Δ''G''⚬ is negative."<ref name="IUPAC Recommendations 1996" /> A strongly exothermic reaction will usually also be exergonic because Δ''H''⚬ makes a major contribution to [[ΔG°|Δ''G''⚬]]. Most of the spectacular chemical reactions that are demonstrated in classrooms are 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==
Examples are numerous: [[combustion]], the [[thermite reaction]], combining strong acids and bases, [[polymerization]]s. As an example in everyday life, [[hand warmer]]s make use of the oxidation of iron to achieve an exothermic reaction:
:4Fe&thinsp; + 3O<sub>2</sub>&thinsp; → 2Fe<sub>2</sub>O<sub>3</sub>&thinsp;                         Δ''H''⚬ = - 1648 kJ/mol

A particularly important class of exothermic reactions is combustion of a hydrocarbon fuel, e.g. the burning of natural gas: 
:CH<sub>4</sub>&thinsp; + 2O<sub>2</sub>&thinsp; → CO<sub>2</sub>&thinsp; + 2H<sub>2</sub>O&thinsp;     Δ''H''⚬ = - 890 kJ/mol
[[File:15. Ослободување на големо количество енергија при согоровуање етанол.webm|thumb|Video of an exothermic reaction.  Ethanol vapor is ignited inside a bottle, causing combustion.]]
These sample reactions are strongly exothermic. 

Uncontrolled exothermic reactions, those leading to [[fire]]s and [[explosion]]s, 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 [[Adenosine triphosphate|ATP]].

==Measurement==
The [[enthalpy]] of a chemical system is essentially its energy. The enthalpy change Δ''H'' 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, Δ{{var|H}}, of a [[combustion]] reaction can be measured particularly accurately. 

The measured heat energy released in an exothermic reaction is converted to Δ''H''⚬ in [[Joule per mole]] (formerly [[Calorie|cal/mol]]). The ''[[Standard state|standard]]'' enthalpy change Δ''H''⚬ is essentially the enthalpy change when the [[Stoichiometry|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, Δ''H''⚬ values are related to [[Bond energy|bond energies]] to a good approximation by:
:Δ{{var|H}}⚬ = total bond energy of reactants − total bond energy of products

[[Image:ac com.svg|300px|thumb|right|An [[energy profile]] of an exothermic reaction]]

In an exothermic reaction, by definition, the enthalpy change has a negative value:
:Δ{{var|H}} = ''H''<sub>products</sub> - ''H''<sub>reactants</sub> < 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:
:2H<sub>2</sub> (g) + O<sub>2</sub> (g) → 2H<sub>2</sub>O (g)
:Δ{{var|H}}⚬ =  −483.6 kJ/mol <ref>{{cite web |url=http://chemistry.osu.edu/~woodward/ch121/ch5_enthalpy.htm |title=Enthalpy (Chapter 5) |access-date=2013-07-20 |url-status=dead |archive-url=https://web.archive.org/web/20130708030319/http://chemistry.osu.edu/~woodward/ch121/ch5_enthalpy.htm |archive-date=2013-07-08 }}</ref>

== See also ==
*[[Chemical thermodynamics]]
*[[Differential scanning calorimetry]]
*[[Endergonic]]
*[[Exergonic]]
*[[Endergonic reaction]]
*[[Exergonic reaction]]
*[[Exothermic process]]
*[[Endothermic reaction]]
*[[Warm-blooded|Endotherm]]

== References ==
{{reflist}}

==External links==

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[[Category:Thermochemistry]]