Editing Chemical decomposition

[[File:Chemical decomposition.gif|thumb|Chemical decomposition]]{{Short description|Breakdown of a chemical species into two or more parts; reverse process of a synthesis reaction}}
'''Chemical decomposition''', or '''chemical breakdown''', is the process or effect of simplifying a single [[molecular entity|chemical entity]] (normal molecule, [[reaction intermediate]], etc.) into two or more fragments.<ref>{{GoldBookRef |title=chemical decomposition |file=C01020 }}</ref> Chemical decomposition is usually regarded and defined as the exact opposite of [[chemical synthesis]]. In short, the chemical reaction in which two or more products are formed from a single reactant is called a decomposition reaction.

The details of a decomposition process are not always well defined. Nevertheless, some activation energy is generally needed to break the involved bonds and as such, higher temperatures generally accelerates decomposition. The net reaction can be an [[endothermic process]], or in the case of spontaneous decompositions, an [[exothermic process]].

The stability of a chemical compound is eventually limited when exposed to extreme environmental conditions such as [[heat]], [[radiation]], [[humidity]], or the [[acidity]] of a [[solvent]]. Because of this chemical decomposition is often an undesired [[chemical reaction]].  However chemical decomposition can be desired, such as in various waste treatment processes.

For example, this method is employed for several analytical techniques, notably [[mass spectrometry]], traditional [[gravimetric analysis]], and [[thermogravimetric analysis]].  Additionally decomposition reactions are used today for a number of other reasons in the production of a wide variety of products. One of these is the explosive breakdown reaction of [[sodium azide]] [(NaN<sub>3</sub>)<sub>2</sub>] into nitrogen gas (N<sub>2</sub>) and sodium (Na). It is this process which powers the life-saving airbags present in virtually all of today's automobiles.<ref>{{Cite web|url=https://prezi.com/1idm8hytmsni/chemical-reactions-in-everyday-life/|title=Chemical reactions in Everyday life|website=prezi.com|language=en|access-date=2017-05-01}}</ref>

Decomposition reactions can be generally classed into three categories; thermal, electrolytic, and photolytic decomposition reactions.<ref>{{Cite web|url=https://amrita.olabs.edu.in/?sub=73&brch=3&sim=80&cnt=1|title=Decomposition Reactions}}</ref>

==Reaction formula==
In the breakdown of a compound into its constituent parts, the generalized reaction for chemical decomposition is:

: AB → A + B (AB represents the reactant that begins the reaction, and A and B represent the products of the reaction)

An example is the [[electrolysis]] of [[water (molecule)|water]] to the gases [[hydrogen]] and [[oxygen]]:

: 2 H<sub>2</sub>O({{serif|l}}) → 2 H<sub>2</sub>({{serif|g}}) + O<sub>2</sub>({{serif|g}})

===Additional examples===
[[File:13. Каталитичко разложување на водород пероксид.webm|thumb|right|280px|An experiment describing catalytic decomposition of hydrogen peroxide, with {{Chem2|[[MnO2]]}} as catalyst. A concentrated hydrogen peroxide solution can be easily decomposed to water and oxygen.]]
An example of a spontaneous (''without'' addition of an external energy source) decomposition is that of [[hydrogen peroxide]] which slowly decomposes into water and oxygen <small>(see video at right</small>):

: 2 H<sub>2</sub>O<sub>2</sub> → 2 H<sub>2</sub>O + O<sub>2</sub>
This reaction is one of the exceptions to the endothermic nature of decomposition reactions.

Other reactions involving decomposition do require the input of external energy. This energy can be in the form of heat, radiation, electricity, or light. The latter being the reason some chemical compounds, such as many prescription medicines, are kept and stored in dark bottles which reduce or eliminate the possibility of light reaching them and initiating decomposition.

When heated, [[carbonate]]s will decompose. A notable exception is [[carbonic acid]], (H<sub>2</sub>CO<sub>3</sub>).<ref>{{Cite web|url=https://ibburke.wordpress.com/2011/03/27/decomposition-of-carbonic-acid/|title=Decomposition of Carbonic Acid Culminating by Elizabeth Burke|last=ibburke|date=2011-03-27|website=ibburke|access-date=2017-03-04}}</ref> Commonly seen as the "fizz" in carbonated beverages, carbonic acid will spontaneously decompose over time into [[carbon dioxide]] and water. The reaction is written as:

: H<sub>2</sub>CO<sub>3</sub> → H<sub>2</sub>O + CO<sub>2</sub>

Other carbonates will decompose when heated to produce their corresponding [[metal]] [[oxide]] and carbon dioxide.<ref>{{Cite web|url=https://Quizlet.com//|title=Synthesis and Decomposition Reactions|last=Walker|first=MS|year=2016|website=Quizlet.com/MSWalker22|series=Online Series in Organic Chemistry|language=en|type=Audio-Video Online Lecture|orig-year=Available now|access-date=2017-03-04|url-access=registration }}</ref> The following equation is an example, where ''M'' represents the given metal:

: ''M''CO<sub>3</sub> → ''M''O + CO<sub>2</sub>

A specific example is that involving [[calcium carbonate]]:
: CaCO<sub>3</sub> → CaO + CO<sub>2</sub>

Metal [[chlorate]]s also decompose when heated. In this type of decomposition reaction, a metal [[chloride]] and oxygen gas are the products. Here, again, ''M'' represents the metal:

: 2 ''M''ClO<sub>3</sub> → 2 ''M''Cl+ 3 O<sub>2</sub>

A common decomposition of a chlorate is in the reaction of [[potassium chlorate]] where oxygen is the product. This can be written as:

: 2 KClO<sub>3</sub> → 2 KCl + 3 O<sub>2</sub>

==See also==
* [[Analytical chemistry]]
* [[Thermal decomposition]]

==References==
{{reflist}}

==External links==
{{wiktionary}}
{{commons category}}
* https://quizlet.com/42968634/types-of-decomposition-reactions-flash-cards/  PDF
* [https://umbbd.ethz.ch Biodegradation database]

{{Authority control}}

[[Category:Inorganic chemistry]]
[[Category:Organic chemistry]]
[[Category:Chemical reactions]]