Editing Chemical reaction (section)

==Reaction types==

===Four basic types===
[[File:Chemical reactions.svg|thumb|500px|Representation of four basic chemical reactions types: synthesis, decomposition, single replacement and double replacement.]]

====Synthesis====
{{Main|Synthesis reaction}}
In a synthesis reaction, two or more simple substances combine to form a more complex substance. These reactions are in the general form:
<chem display="block">A + B->AB</chem>

Two or more reactants yielding one product is another way to identify a synthesis reaction. One example of a synthesis reaction is the combination of [[iron]] and [[sulfur]] to form [[iron(II) sulfide]]:
<chem display="block">8Fe + S8->8FeS</chem>

Another example is simple hydrogen gas combined with simple oxygen gas to produce a more complex substance, such as water.<ref name="to react or not to react">[http://utahscience.oremjr.alpine.k12.ut.us/sciber99/8th/matter/sciber/chemtype.htm To react or not to react?] {{Webarchive|url=https://web.archive.org/web/20150110214558/http://utahscience.oremjr.alpine.k12.ut.us/sciber99/8th/matter/sciber/chemtype.htm |date=2015-01-10 }} Utah State Office of Education. Retrieved 4 June 2011.</ref>

====Decomposition====
{{Main|Decomposition reaction}}
A decomposition reaction is when a more complex substance breaks down into its more simple parts. It is thus the opposite of a synthesis reaction and can be written as<ref name="to react or not to react"/>
<chem display="block">AB->A + B</chem>

One example of a decomposition reaction is the [[electrolysis]] of water to make [[oxygen]] and [[hydrogen]] gas:
<chem display="block">2H2O->2H2 + O2</chem>

====Single displacement====
In a [[single displacement reaction]], a single uncombined element replaces another in a compound; in other words, one element trades places with another element in a compound<ref name="to react or not to react" /> These reactions come in the general form of:
<chem display="block">A + BC->AC + B</chem>

One example of a single displacement reaction is when [[magnesium]] replaces hydrogen in water to make solid [[magnesium hydroxide]] and hydrogen gas:
<chem display="block">Mg + 2H2O->Mg(OH)2 (v) + H2 (^)</chem>

====Double displacement====
In a [[double replacement reaction|double displacement reaction]], the anions and cations of two compounds switch places and form two entirely different compounds. These reactions are in the general form:<ref name="to react or not to react" />
<chem display="block">AB + CD->AD + CB</chem>

For example, when [[barium chloride]] (BaCl<sub>2</sub>) and [[magnesium sulfate]] (MgSO<sub>4</sub>) react, the SO<sub>4</sub><sup>2−</sup> anion switches places with the 2Cl<sup>−</sup> anion, giving the compounds BaSO<sub>4</sub> and MgCl<sub>2</sub>.

Another example of a double displacement reaction is the reaction of [[lead(II) nitrate]] with [[potassium iodide]] to form [[lead(II) iodide]] and [[potassium nitrate]]:
<chem display="block">Pb(NO3)2 + 2KI->PbI2(v) + 2KNO3</chem>

===Forward and backward reactions===
According to [[Le Chatelier's Principle]], reactions may proceed in the forward or reverse direction until they end or reach [[Equilibrium chemistry|equilibrium]].<ref name="libretext1">{{Cite web |date=2016-08-05 |title=8.3: Le Châtelier's Principle |url=https://chem.libretexts.org/Courses/University_of_Kentucky/UK%3A_CHE_103_-_Chemistry_for_Allied_Health_(Soult)/Chapters/Chapter_8%3A_Properties_of_Solutions/8.3%3A_Le_Ch%C3%A2telier%27s_Principle |access-date=2023-04-11 |website=Chemistry LibreTexts |language=en}}</ref>

====Forward reactions====
Reactions that proceed in the forward direction (from left to right) to approach equilibrium are often called [[spontaneous reaction]]s, that is, <math>\Delta G</math> is negative, which means that if they occur at constant temperature and pressure, they decrease the [[Gibbs free energy]] of the reaction. They require less energy to proceed in the forward direction.<ref name="libretext2">{{Cite web |date=2016-08-05 |title=11.5: Spontaneous Reactions and Free Energy |url=https://chem.libretexts.org/Courses/University_of_Kentucky/UK%3A_CHE_103_-_Chemistry_for_Allied_Health_(Soult)/Chapters/Chapter_11%3A_Properties_of_Reactions/11.5%3A_Spontaneous_Reactions_and_Free_Energy#:~:text=The%20forward,form. |access-date=2023-04-11 |website=Chemistry LibreTexts |language=en}}</ref> Reactions are usually written as forward reactions in the direction in which they are spontaneous. Examples:
* Reaction of hydrogen and oxygen to form water.
:{{chem|2H|2}} + {{chem|O|2}} {{eqm}} {{chem|2H|2|O}}
* Dissociation of [[acetic acid]] in water into [[acetate]] ions and [[hydronium ion]]s.
:{{chem|CH|3|COOH}} + {{chem|H|2|O}} {{eqm}} {{chem|CH|3|COO|-}} + {{chem|H|3|O|+}}

====Backward reactions====
Reactions that proceed in the backward direction to approach equilibrium are often called [[non-spontaneous reaction]]s, that is, <math>\Delta G</math> is positive, which means that if they occur at constant temperature and pressure, they increase the [[Gibbs free energy]] of the reaction. They require input of energy to proceed in the forward direction.<ref name="libretext2" /><ref name="libretext3">{{Cite web |date=2016-06-27 |title=20.3: Spontaneous and Nonspontaneous Reactions |url=https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Introductory_Chemistry_(CK-12)/20%3A_Entropy_and_Free_Energy/20.03%3A_Spontaneous_and_Nonspontaneous_Reactions#:~:text=In%20the%20reverse,spontaneous. |access-date=2023-04-11 |website=Chemistry LibreTexts |language=en}}</ref> Examples include:
* Charging a normal DC battery (consisting of [[electrolytic cell]]s) from an external electrical power source<ref>{{Cite web |date=2013-10-02 |title=Electrolytic Cells |url=https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Electrochemistry/Electrolytic_Cells |access-date=2023-04-11 |website=Chemistry LibreTexts |language=en}}</ref>
* [[Photosynthesis]] driven by absorption of [[electromagnetic radiation]] usually in the form of sunlight<ref>{{Cite web |date=2016-05-26 |title=Photosynthesis of Exoplanet Plants |url=https://chem.libretexts.org/Ancillary_Materials/Exemplars_and_Case_Studies/Exemplars/Physics_and_Astronomy/Photosynthesis_of_Exoplanet_Plants |access-date=2023-04-11 |website=Chemistry LibreTexts |language=en}}</ref>

:{{underset|carbon<br />dioxide|CO<sub>2</sub>}} + {{underset| water |H<sub>2</sub>O}} + {{underset|light energy|photons}} → {{underset|carbohydrate|[CH<sub>2</sub>O]}} + {{underset| oxygen |O<sub>2</sub>}}

===Combustion===
In a [[combustion]] reaction, an element or compound reacts with an oxidant, usually [[oxygen]], often producing energy in the form of [[heat]] or [[light]]. Combustion reactions frequently involve a [[hydrocarbon]]. For instance, the combustion of 1 mole (114 g) of octane in oxygen
<chem display="block">C8H18(l) + 25/2 O2(g)->8CO2 + 9H2O(l)</chem>

releases 5500 kJ. A combustion reaction can also result from [[carbon]], [[magnesium]] or [[sulfur]] reacting with oxygen.<ref>{{cite book |author1=Wilbraham, Antony |author2= Stanley, Dennis|author3=Waterman, Edward|author4=Matta, Michael|title=Chemistry |date=2012 |publisher=Pearson |isbn=9780132525763|pages=734–735 }}</ref>
<chem display="block">2Mg(s) + O2->2MgO(s)</chem>
<chem display="block">S(s) + O2(g)->SO2(g)</chem>

===Oxidation and reduction===
[[File:redox reaction.png|thumb|right|250px|Illustration of a redox reaction]]
[[File:Common-salt.jpg|thumb|right|250px|[[Sodium chloride]] is formed through the redox reaction of sodium metal and chlorine gas]]

[[Redox]] reactions can be understood in terms of the transfer of electrons from one involved species ([[reducing agent]]) to another ([[oxidizing agent]]). In this process, the former species is ''oxidized'' and the latter is ''reduced''. Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation is better defined as an increase in [[oxidation state]] of atoms and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always change the oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving [[covalent]] bonds).<ref>{{cite encyclopedia | author = Glusker, Jenny P. |author-link=Jenny Glusker| contribution = Structural Aspects of Metal Liganding to Functional Groups in Proteins | editor = Christian B. Anfinsen | url = https://books.google.com/books?id=HvARsi6S-b0C&pg=PA7 | title = Advances in Protein Chemistry | volume = 42 | publisher = [[Academic Press]] | location = San Diego | year = 1991 | isbn = 978-0-12-034242-6 | page = 7}}</ref><ref>{{ cite encyclopedia | author = Guo, Liang-Hong | author2 = Allen, H. | author3 = Hill, O. | contribution = Direct Electrochemistry of Proteins and Enzymes | editor = A.G. Sykes | url = https://books.google.com/books?id=qnRkjATn0YUC&pg=PA359 | title = Advances in Inorganic Chemistry | volume = 36 | publisher = [[Academic Press]] | location = San Diego | year = 1991 | isbn = 978-0-12-023636-7 | page = 359}}</ref>

In the following redox reaction, hazardous [[sodium]] metal reacts with toxic [[chlorine]] gas to form the ionic compound [[sodium chloride]], or common table salt:
<chem display="block">2Na(s) + Cl2(g)->2NaCl(s)</chem>

In the reaction, sodium metal goes from an oxidation state of 0 (a pure element) to +1: in other words, the sodium lost one electron and is said to have been oxidized. On the other hand, the chlorine gas goes from an oxidation of 0 (also a pure element) to −1: the chlorine gains one electron and is said to have been reduced. Because the chlorine is the one reduced, it is considered the electron acceptor, or in other words, induces oxidation in the sodium – thus the chlorine gas is considered the oxidizing agent. Conversely, the sodium is oxidized or is the electron donor, and thus induces a reduction in the other species and is considered the ''reducing agent''.

Which of the involved reactants would be a reducing or oxidizing agent can be predicted from the [[electronegativity]] of their elements. Elements with low electronegativities, such as most metals, easily donate electrons and oxidize – they are reducing agents. On the contrary, many oxides or ions with high oxidation numbers of their non-oxygen atoms, such as {{chem|link=hydrogen peroxide|H|2|O|2}}, {{chem|link=permanganate|MnO|4|-}}, {{chem|link=chromium trioxide|CrO|3}}, {{chem|link=dichromate|Cr|2|O|7|2-}}, or {{chem|link=Osmium(VIII) oxide|OsO|4}}, can gain one or two extra electrons and are strong oxidizing agents.

For some [[main-group element]]s the number of electrons donated or accepted in a redox reaction can be predicted from the [[electron configuration]] of the reactant element. Elements try to reach the low-energy [[noble gas]] configuration, and therefore alkali metals and halogens will donate and accept one electron, respectively. Noble gases themselves are chemically inactive.<ref>[[#Wiberg|Wiberg]], pp. 289–290</ref>

The overall redox reaction [[Electrochemistry#Balancing redox reactions|can be balanced]] by combining the oxidation and reduction half-reactions multiplied by coefficients such that the number of electrons lost in the oxidation equals the number of electrons gained in the reduction.

An important class of redox reactions are the electrolytic [[Electrochemistry|electrochemical]] reactions, where electrons from the power supply at the negative electrode are used as the reducing agent and electron withdrawal at the positive electrode as the oxidizing agent. These reactions are particularly important for the production of chemical elements, such as [[chlorine]]<ref>[[#Wiberg|Wiberg]], p. 409</ref> or [[aluminium]]. The reverse process, in which electrons are released in redox reactions and chemical energy is converted to electrical energy, is possible and used in [[Electric battery|batteries]].

===Complexation===
[[File:Ferrocene-from-xtal-3D-balls.png|160px|thumb|[[Ferrocene]] – an iron atom sandwiched between two C<sub>5</sub>H<sub>5</sub> [[ligand]]s]]
In complexation reactions, several [[ligand]]s react with a metal atom to form a [[coordination complex]]. This is achieved by providing [[lone pair]]s of the ligand into empty [[Atomic orbital|orbitals]] of the metal atom and forming [[dipolar bond]]s. The ligands are [[Lewis base]]s, they can be both ions and neutral molecules, such as carbon monoxide, ammonia or water. The number of ligands that react with a central metal atom can be found using the [[18-electron rule]], saying that the [[valence shell]]s of a [[transition metal]] will collectively accommodate 18 [[electron]]s, whereas the symmetry of the resulting complex can be predicted with the [[crystal field theory]] and [[ligand field theory]]. Complexation reactions also include [[ligand exchange]], in which one or more ligands are replaced by another, and redox processes which change the oxidation state of the central metal atom.<ref>[[#Wiberg|Wiberg]], pp. 1180–1205</ref>

===Acid–base reactions===
In the [[Brønsted–Lowry acid–base theory]], an [[acid–base reaction]] involves a transfer of [[proton]]s (H<sup>+</sup>) from one species (the [[acid]]) to another (the [[base (chemistry)|base]]). When a proton is removed from an acid, the resulting species is termed that acid's [[conjugate acid|conjugate base]]. When the proton is accepted by a base, the resulting species is termed that base's [[conjugate acid]].<ref>{{GoldBookRef|title=conjugate acid–base pair|file=C01266}}</ref> In other words, acids act as proton donors and bases act as proton acceptors according to the following equation:
<chem display="block">\underset{acid}{HA} + \underset{base}{B} <=> \underset{conjugated\ base}{A^-} + \underset{conjugated\ acid}{HB+}</chem>

The reverse reaction is possible, and thus the acid/base and conjugated base/acid are always in equilibrium. The equilibrium is determined by the [[acid dissociation constant|acid and base dissociation constants]] (''K''<sub>a</sub> and ''K''<sub>b</sub>) of the involved substances. A special case of the acid-base reaction is the [[neutralization (chemistry)|neutralization]] where an acid and a base, taken at the exact same amounts, form a neutral [[Salt (chemistry)|salt]].

Acid-base reactions can have different definitions depending on the acid-base concept employed. Some of the most common are:
* [[Acid–base reaction#Arrhenius definition|Arrhenius]] definition: Acids dissociate in water releasing H<sub>3</sub>O<sup>+</sup> ions; bases dissociate in water releasing OH<sup>−</sup> ions.
* [[Brønsted–Lowry acid–base theory|Brønsted–Lowry]] definition: Acids are proton (H<sup>+</sup>) donors, bases are proton acceptors; this includes the Arrhenius definition.
* [[Acid–base reaction#Lewis definition|Lewis]] definition: Acids are electron-pair acceptors, and bases are electron-pair donors; this includes the Brønsted-Lowry definition.

===Precipitation===
[[File:Chemical precipitation diagram multilang.svg|thumb|Precipitation]]
[[Precipitation (chemistry)|Precipitation]] is the formation of a solid in a solution or inside another solid during a chemical reaction. It usually takes place when the concentration of dissolved ions exceeds the [[solubility]] limit<ref>{{GoldBookRef|title=precipitation|file=P04795}}</ref> and forms an insoluble salt. This process can be assisted by adding a precipitating agent or by the removal of the solvent. Rapid precipitation results in an [[amorphous]] or microcrystalline residue and a slow process can yield single [[crystal]]s. The latter can also be obtained by [[Recrystallization (chemistry)|recrystallization]] from microcrystalline salts.<ref>{{cite encyclopedia | author = Wingender, Jörg | author2 = Ortanderl, Stefanie | contribution = Ausfällung | title = Römpp Chemie-Lexikon | publisher = [[Thieme Medical Publishers|Thieme]] | date = July 2009}}</ref>

===Solid-state reactions===
Reactions can take place between two solids. However, because of the relatively small [[diffusion]] rates in solids, the corresponding chemical reactions are very slow in comparison to liquid and gas phase reactions. They are accelerated by increasing the reaction temperature and finely dividing the reactant to increase the contacting surface area.<ref>{{cite encyclopedia | editor = Erwin Riedel | author = Meyer, H. Jürgen | title = Modern Inorganic Chemistry | contribution = Festkörperchemie | edition = 3rd | url = https://books.google.com/books?id=HwY4be5bH_sC&pg=PA171 | language = de | publisher = [[Walter de Gruyter|de Gruyter]] | year = 2007 | isbn = 978-3-11-019060-1 | page = 171}}</ref>

===Reactions at the solid/gas interface===
The reaction can take place at the solid|gas interface, surfaces at very low pressure such as [[ultra-high vacuum]]. Via [[scanning tunneling microscopy]], it is possible to observe reactions at the solid|gas interface in real space, if the time scale of the reaction is in the correct range.<ref>{{Cite journal | doi = 10.1126/science.278.5345.1931| title = Atomic and Macroscopic Reaction Rates of a Surface-Catalyzed Reaction| journal = Science| volume = 278| issue = 5345| pages = 1931–4| year = 1997| last1 = Wintterlin | first1 = J.| pmid=9395392| bibcode = 1997Sci...278.1931W}}</ref><ref>{{Cite journal | doi = 10.1021/ja302593v| title = Oxidation of an Organic Adlayer: A Bird's Eye View| journal = Journal of the American Chemical Society| volume = 134| issue = 21| pages = 8817–8822| year = 2012| last1 = Waldmann | first1 = T. | last2 = Künzel | first2 = D. | last3 = Hoster | first3 = H.E. | last4 = Groß | first4 = A. | last5 = Behm | first5 = R.J.R. | pmid=22571820| bibcode = 2012JAChS.134.8817W}}</ref> Reactions at the solid|gas interface are in some cases related to catalysis.

===Photochemical reactions===
[[File:Paterno-Buchi reaction.svg|thumb|In this [[Paterno–Büchi reaction]], a photoexcited carbonyl group is added to an unexcited [[olefin]], yielding an [[oxetane]].]]
In [[Photochemistry|photochemical reactions]], atoms and molecules absorb energy ([[photon]]s) of the illumination light and convert it into an [[excited state]]. They can then release this energy by breaking chemical bonds, thereby producing radicals. Photochemical reactions include hydrogen–oxygen reactions, [[radical polymerization]], [[chain reaction]]s and [[rearrangement reaction]]s.<ref>[[#Atkins|Atkins]], pp. 937–950</ref>

Many important processes involve photochemistry. The premier example is [[photosynthesis]], in which most plants use solar energy to convert [[carbon dioxide]] and water into [[glucose]], disposing of [[oxygen]] as a side-product. Humans rely on photochemistry for the formation of vitamin D, and [[visual perception|vision]] is initiated by a photochemical reaction of [[rhodopsin]].<ref name=rh>{{cite encyclopedia | author = Kandori, Hideki | contribution = Retinal Binding Proteins | editor = Dugave, Christophe | year = 2006 | url = https://books.google.com/books?id=udSCHPq5Ii0C&pg=PA56 | title = Cis-trans Isomerization in Biochemistry | publisher = [[John Wiley & Sons|Wiley-VCH]] | page = 56 | isbn = 978-3-527-31304-4}}</ref> In [[fireflies]], an [[enzyme]] in the abdomen catalyzes a reaction that results in [[bioluminescence]].<ref>{{cite book | author = Saunders, David Stanley | year = 2002 | url = https://books.google.com/books?id=3qJOw5Gh_UMC&pg=PA179 | title = Insect clocks | edition = Third | publisher = [[Elsevier]] | location = Amsterdam | page = 179 | isbn = 978-0-444-50407-4}}</ref> Many significant photochemical reactions, such as ozone formation, occur in the Earth atmosphere and constitute [[atmospheric chemistry]].