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{{short description|Process that results in the interconversion of chemical species}}
{{pp|small=yes}}
[[File:ThermiteReaction.jpg|thumb|A [[thermite]] reaction using iron(III) oxide. The sparks flying outwards are globules of molten iron trailing smoke in their wake.]]

A '''chemical reaction''' is a process that leads to the [[chemistry|chemical]] transformation of one set of [[chemical substance]]s to another.<ref>{{GoldBookRef|title=chemical reaction|file=C01033}}</ref> When chemical reactions occur, the [[atom]]s are rearranged and the reaction is accompanied by an [[Gibbs free energy|energy change]] as new products are generated. Classically, [[chemical]] reactions encompass changes that only involve the positions of [[electron]]s in the forming and breaking of [[chemical bond]]s between [[atom]]s, with no change to the [[Atomic nucleus|nuclei]] (no change to the elements present), and can often be described by a [[chemical equation]]. [[Nuclear chemistry]] is a sub-discipline of chemistry that involves the chemical reactions of [[unstable]] and [[radioactive]] [[Chemical element|element]]s where both electronic and nuclear changes can occur.

The substance (or substances) initially involved in a chemical reaction are called [[reagent|reactants or reagents]]. Chemical reactions are usually characterized by a [[chemical change]], and they yield one or more [[Product (chemistry)|products]], which usually have properties different from the reactants. Reactions often consist of a sequence of individual sub-steps, the so-called [[elementary reaction]]s, and the information on the precise course of action is part of the [[reaction mechanism]]. Chemical reactions are described with [[chemical equation]]s, which symbolically present the starting materials, end products, and sometimes intermediate products and reaction conditions.

Chemical reactions happen at a characteristic [[reaction rate]] at a given temperature and chemical concentration. Some reactions produce [[heat]] and are called [[exothermic reaction]]s, while others may require heat to enable the reaction to occur, which are called [[Endothermic process|endothermic reactions]]. Typically, reaction rates increase with increasing temperature because there is more [[thermal energy]] available to reach the activation energy necessary for breaking bonds between atoms.

A reaction may be classified as [[redox]] in which [[oxidation]] and [[Redox|reduction]] occur or non-redox in which there is no oxidation and reduction occurring. Most simple redox reactions may be classified as a combination, decomposition, or single displacement reaction.

Different chemical reactions are used during [[chemical synthesis]] in order to obtain the desired product. In [[biochemistry]], a consecutive series of chemical reactions (where the product of one reaction is the reactant of the next reaction) form [[metabolic pathway]]s. These reactions are often [[catalysis|catalyzed]] by protein [[enzyme]]s. Enzymes increase the rates of biochemical reactions, so that [[metabolism|metabolic]] syntheses and decompositions impossible under ordinary conditions can occur at the temperature and concentrations present within a [[Cell (biology)|cell]].

The general concept of a chemical reaction has been extended to reactions between entities smaller than atoms, including [[nuclear reactions]], [[radioactive decay]]s and reactions between [[elementary particle]]s, as described by [[quantum field theory]].

==History==
[[File:David - Portrait of Monsieur Lavoisier (cropped).jpg|thumb|[[Antoine Lavoisier]] developed the theory of combustion as a chemical reaction with oxygen.]]

Chemical reactions such as combustion in fire, [[fermentation (biochemistry)|fermentation]] and the reduction of ores to metals were known since antiquity. Initial theories of transformation of materials were developed by Greek philosophers, such as the [[Classical element|Four-Element Theory]] of [[Empedocles]] stating that any substance is composed of the four basic elements – fire, water, air and earth. In the [[Middle Ages]], chemical transformations were studied by [[alchemist]]s. They attempted, in particular, to convert [[lead]] into [[gold]], for which purpose they used reactions of lead and lead-copper alloys with [[sulfur]].<ref>{{Cite journal | doi = 10.1002/ciuz.19730070604| title = Neuere Interpretationsmöglichkeiten der Alchemie| journal = Chemie in unserer Zeit| volume = 7| issue = 6| pages = 177–181| year = 1973| last1 = Weyer | first1 = J. }}</ref>

The artificial production of chemical substances already was a central goal for medieval alchemists.<ref>See {{Cite book|last=Newman|first=William R.|author-link=William R. Newman|year=2004|title=Promethean Ambitions: Alchemy and the Quest to Perfect Nature|location=Chicago|publisher=University of Chicago Press|isbn=9780226575247}}</ref> Examples include the synthesis of [[ammonium chloride]] from [[Organic compound|organic substances]] as described in the works (c. 850–950) attributed to [[Jabir ibn Hayyan|Jābir ibn Ḥayyān]],<ref>{{Cite book|last=Kraus|first=Paul|author-link=Paul Kraus (Arabist)|year=1942–1943|title=Jâbir ibn Hayyân: Contribution à l'histoire des idées scientifiques dans l'Islam. I. Le corpus des écrits jâbiriens. II. Jâbir et la science grecque|publisher=[[Institut Français d'Archéologie Orientale]]|location=Cairo|oclc=468740510|isbn=9783487091150}}, vol. II, pp. 41–42.</ref> or the production of [[mineral acids]] such as [[sulfuric acid|sulfuric]] and [[nitric acid]]s by later alchemists, starting from c. 1300.<ref>{{Cite journal|last1=Karpenko|first1=Vladimír|last2=Norris|first2=John A.|year=2002|title=Vitriol in the History of Chemistry|journal=Chemické listy|volume=96|issue=12|pages=997–1005|url=http://www.chemicke-listy.cz/ojs3/index.php/chemicke-listy/article/view/2266}}</ref> The production of mineral acids involved the heating of sulfate and nitrate minerals such as [[copper sulfate]], [[alum]] and [[Potassium nitrate|saltpeter]]. In the 17th century, [[Johann Rudolph Glauber]] produced [[hydrochloric acid]] and [[sodium sulfate]] by reacting sulfuric acid and [[sodium chloride]]. With the development of the [[lead chamber process]] in 1746 and the [[Leblanc process]], allowing large-scale production of sulfuric acid and [[sodium carbonate]], respectively, chemical reactions became implemented into the industry. Further optimization of sulfuric acid technology resulted in the [[contact process]] in the 1880s,<ref>{{cite book | author = Friedman, Leonard J. | author2 = Friedman, Samantha J. | url = http://www.aiche-cf.org/Clearwater/2008/Paper2/8.2.7.pdf | title = The History of the Contact Sulfuric Acid Process | year = 2008 | publisher = Acid Engineering & Consulting, Inc. | location = Boca Raton, Florida}}</ref> and the [[Haber process]] was developed in 1909–1910 for [[ammonia]] synthesis.<ref>{{cite encyclopedia | editor = Lesch, John E. | url = https://books.google.com/books?id=VJIztvolC8cC&pg=PA170 | title = The German Chemical Industry in the Twentieth Century | author = Stranges, Anthony N. | contribution = Germany's synthetic fuel industry, 1935–1940 | publisher = [[Kluwer Academic Publishers]] | year = 2000 | isbn = 978-0-7923-6487-0 | page = 170}}</ref>

From the 16th century, researchers including [[Jan Baptist van Helmont]], [[Robert Boyle]], and [[Isaac Newton]] tried to establish theories of experimentally observed chemical transformations. The [[phlogiston theory]] was proposed in 1667 by [[J. J. Becher|Johann Joachim Becher]]. It postulated the existence of a fire-like element called "phlogiston", which was contained within combustible bodies and released during [[combustion]]. This proved to be false in 1785 by [[Antoine Lavoisier]] who found the correct explanation of the combustion as a reaction with oxygen from the air.<ref>[[#Brock|Brock]], pp. 34–55</ref>

[[Joseph Louis Gay-Lussac]] recognized in 1808 that gases always react in a certain relationship with each other. Based on this idea and the atomic theory of [[John Dalton]], [[Joseph Proust]] had developed the [[law of definite proportions]], which later resulted in the concepts of [[stoichiometry]] and [[chemical equation]]s.<ref>[[#Brock|Brock]], pp. 104–107</ref>

Regarding the [[organic chemistry]], it was long believed that compounds obtained from living organisms were too complex to be obtained [[Chemical synthesis|synthetically]]. According to the concept of [[vitalism]], organic matter was endowed with a "vital force" and distinguished from inorganic materials. This separation was ended however by the synthesis of [[urea]] from inorganic precursors by [[Friedrich Wöhler]] in 1828. Other chemists who brought major contributions to organic chemistry include [[Alexander William Williamson]] with his [[Williamson ether synthesis|synthesis]] of [[ether]]s and [[Christopher Kelk Ingold]], who, among many discoveries, established the mechanisms of [[substitution reaction]]s.

== Characteristics ==
{{Expand section|date=November 2020}}
The general characteristics of chemical reactions are:
* [[Gas evolution reaction|Evolution of a gas]]
* Formation of a [[Precipitation (chemistry)|precipitate]]
* Change in [[temperature]]
* Change in [[State of matter|state]]

==Equations==
[[File:Combustion reaction of methane.jpg|thumb|450px|As seen from the equation {{chem2|CH4 + 2O2 -> CO2 + 2 H2O}}, a coefficient of 2 must be placed before the [[oxygen]] gas on the reactants side and before the [[properties of water|water]] on the products side in order for, as per the law of conservation of mass, the quantity of each element does not change during the reaction.]]

{{main|Chemical equation}}

[[Chemical equation]]s are used to graphically illustrate chemical reactions. They consist of [[chemical formula|chemical]] or [[structural formula]]s of the reactants on the left and those of the products on the right. They are separated by an arrow (→) which indicates the direction and type of the reaction; the arrow is read as the word "yields".<ref>{{cite book|author=Myers, Richard |title=The Basics of Chemistry|year=2009|url=https://books.google.com/books?id=oS50J3-IfZsC |publisher=[[Greenwood Publishing Group]]|isbn=978-0-313-31664-7|page=55}}</ref> The tip of the arrow points in the direction in which the reaction proceeds. A double arrow ({{eqm}}) pointing in opposite directions is used for [[Chemical equilibrium|equilibrium reactions]]. Equations should be balanced according to the [[stoichiometry]], the number of atoms of each species should be the same on both sides of the equation. This is achieved by scaling the number of involved molecules (A, B, C and D in a schematic example below) by the appropriate integers ''a, b, c'' and ''d''.<ref>{{GoldBookRef|title=chemical reaction equation|file=C01034}}</ref>

:{{chem2|''a'' A + ''b'' B -> ''c'' C + ''d'' D}}

More elaborate reactions are represented by reaction schemes, which in addition to starting materials and products show important intermediates or [[transition state]]s. Also, some relatively minor additions to the reaction can be indicated above the reaction arrow; examples of such additions are water, heat, illumination, a [[catalyst]], etc. Similarly, some minor products can be placed below the arrow, often with a minus sign.
[[File:Baeyer-Villiger-Oxidation-V1.svg|thumb|center|750px|An example of organic reaction: [[Baeyer–Villiger oxidation|oxidation]] of [[ketone]]s to [[ester]]s with a [[Peroxy acid|peroxycarboxylic acid]]]]
{{clear}}
[[Retrosynthetic analysis]] can be applied to design a complex synthesis reaction. Here the analysis starts from the products, for example by splitting selected chemical bonds, to arrive at plausible initial reagents. A special arrow (⇒) is used in retro reactions.<ref>{{Cite journal | doi = 10.1039/CS9881700111| title = Robert Robinson Lecture. Retrosynthetic thinking?essentials and examples| journal = Chemical Society Reviews| volume = 17| pages = 111–133| year = 1988| last1 = Corey | first1 = E.J.}}</ref>

==Elementary reactions==
The [[elementary reaction]] is the smallest division into which a chemical reaction can be decomposed, it has no intermediate products.<ref>{{GoldBookRef|title=elementary reaction|file=E02035}}</ref> Most experimentally observed reactions are built up from many elementary reactions that occur in parallel or sequentially. The actual sequence of the individual elementary reactions is known as [[reaction mechanism]]. An elementary reaction involves a few molecules, usually one or two, because of the low probability for several molecules to meet at a certain time.<ref>{{cite encyclopedia | author = Frenking, Gernot | contribution = Elementarreaktionen | title = Römpp Chemie-Lexikon | publisher = [[Thieme Medical Publishers|Thieme]] | year = 2006}}</ref>

[[File:Azobenzene isomerization de.svg|thumb|300px|Isomerization of [[azobenzene]], induced by light (hν) or heat (Δ)]]
The most important elementary reactions are unimolecular and bimolecular reactions. Only one molecule is involved in a unimolecular reaction; it is transformed by isomerization or a [[dissociation (chemistry)|dissociation]] into one or more other molecules. Such reactions require the addition of energy in the form of heat or light. A typical example of a unimolecular reaction is the [[cis–trans isomerism|cis–trans]] [[isomerization]], in which the cis-form of a compound converts to the trans-form or vice versa.<ref name=rh/>

In a typical [[dissociation (chemistry)|dissociation]] reaction, a bond in a molecule splits ('''ruptures''') resulting in two molecular fragments. The splitting can be [[homolysis (chemistry)|homolytic]] or [[Heterolysis (chemistry)|heterolytic]]. In the first case, the bond is divided so that each product retains an electron and becomes a neutral [[Radical (chemistry)|radical]]. In the second case, both electrons of the chemical bond remain with one of the products, resulting in charged [[ion]]s. Dissociation plays an important role in triggering [[chain reaction]]s, such as [[Oxyhydrogen|hydrogen–oxygen]] or [[polymerization]] reactions.

:<chem>AB -> A + B</chem>
: <small> Dissociation of a molecule AB into fragments A and B </small>

For [[bimolecular]] reactions, two molecules collide and react with each other. Their merger is called [[chemical synthesis]] or an [[addition reaction]].
:<chem>A + B -> AB</chem>

Another possibility is that only a portion of one molecule is transferred to the other molecule. This type of reaction occurs, for example, in [[redox]] and acid-base reactions. In redox reactions, the transferred particle is an electron, whereas in acid-base reactions it is a proton. This type of reaction is also called [[Salt metathesis reaction|metathesis]].
:<chem>HA + B -> A + HB</chem>
for example
:<chem>NaCl + AgNO3 -> NaNO3 + AgCl(v)</chem>

==Chemical equilibrium==
{{Main|Chemical equilibrium}}
Most chemical reactions are reversible; that is, they can and do run in both directions. The forward and reverse reactions are competing with each other and differ in [[Chemical kinetics|reaction rates]]. These rates depend on the concentration and therefore change with the time of the reaction: the reverse rate gradually increases and becomes equal to the rate of the forward reaction, establishing the so-called chemical equilibrium. The time to reach equilibrium depends on parameters such as temperature, pressure, and the materials involved, and is determined by the [[Principle of minimum energy|minimum free energy]]. In equilibrium, the [[Gibbs free energy]] of reaction must be zero. The pressure dependence can be explained with the [[Le Chatelier's principle]]. For example, an increase in pressure due to decreasing volume causes the reaction to shift to the side with fewer moles of gas.<ref>[[#Atkins|Atkins]], p. 114.</ref>

The reaction yield stabilizes at equilibrium but can be increased by removing the product from the reaction mixture or changed by increasing the temperature or pressure. A change in the concentrations of the reactants does not affect the equilibrium constant but does affect the equilibrium position.

==Thermodynamics==
Chemical reactions are determined by the laws of [[thermodynamics]]. Reactions can proceed by themselves if they are [[exergonic]], that is if they release free energy. The associated free energy change of the reaction is composed of the changes of two different thermodynamic quantities, [[enthalpy]] and [[entropy]]:<ref>[[#Atkins|Atkins]], pp. 106–108</ref>

:; <math>\Delta G = \Delta H - T \cdot \Delta S</math>.
:: <small> {{mvar|G}}: free energy, {{mvar|H}}: enthalpy, {{mvar|T}}: temperature, {{mvar|S}}: entropy, {{math|Δ}}: difference (change between original and product) </small>

Reactions can be [[Exothermic reaction|exothermic]], where Δ''H'' is negative and energy is released. Typical examples of exothermic reactions are [[combustion]], [[Precipitation (chemistry)|precipitation]] and [[crystallization]], in which ordered solids are formed from disordered gaseous or liquid phases. In contrast, in [[endothermic]] reactions, heat is consumed from the environment. This can occur by increasing the entropy of the system, often through the formation of gaseous or dissolved reaction products, which have higher entropy. Since the entropy term in the free-energy change increases with temperature, many endothermic reactions preferably take place at high temperatures. On the contrary, many exothermic reactions such as crystallization occur preferably at lower temperatures. A change in temperature can sometimes reverse the sign of the enthalpy of a reaction, as for the [[carbon monoxide]] reduction of [[molybdenum dioxide]]:
:<chem>2CO(g) + MoO2(s) -> 2CO2(g) + Mo(s)</chem>; <math>\Delta H^o = +21.86 \ \text{kJ at 298 K}</math>
This reaction to form [[carbon dioxide]] and [[molybdenum]] is endothermic at low temperatures, becoming less so with increasing temperature.<ref name="ReacWeb">{{Cite web|url=https://www.crct.polymtl.ca/reacweb.htm|title=F*A*C*T - REACTION-Web|website=www.crct.polymtl.ca}}</ref> Δ''H''° is zero at {{val|1855|ul=K}}, and the reaction becomes exothermic above that temperature.

Changes in temperature can also reverse the direction tendency of a reaction. For example, the [[water gas shift reaction]]
:<chem>CO(g) + H2O({v}) <=> CO2(g) + H2(g)</chem>
is favored by low temperatures, but its reverse is favored by high temperatures. The shift in reaction direction tendency occurs at {{val|1100|u=K}}.<ref name="ReacWeb" />

Reactions can also be characterized by their [[internal energy]] change, which takes into account changes in the entropy, volume and [[chemical potential]]s. The latter depends, among other things, on the [[activity (chemistry)|activities]] of the involved substances.<ref>[[#Atkins|Atkins]], p. 150</ref>

:; <math>{d}U = T\cdot {d}S - p\cdot {d}V + \mu\cdot {d}n</math>
:: <small> {{mvar|U}}: internal energy, {{mvar|S}}: entropy, {{mvar|p}}: pressure, {{mvar|μ}}: chemical potential, {{mvar|n}}: number of molecules, {{mvar|d}}: [[differential calculus|small change sign]] </small>

==Kinetics==
The speed at which reactions take place is studied by [[Chemical kinetics|reaction kinetics]]. The rate depends on various parameters, such as:
* [[Reactant]] concentrations, which usually make the reaction happen at a faster rate if raised through increased collisions per unit of time. Some reactions, however, have rates that are ''independent'' of reactant concentrations, due to a limited number of catalytic sites.  These are called [[Rate law#Zero-order reactions|zero order reactions]].
* [[Surface area]] available for contact between the reactants, in particular solid ones in heterogeneous systems. Larger surface areas lead to higher reaction rates.
* [[Pressure]] – increasing the pressure decreases the volume between molecules and therefore increases the frequency of collisions between the molecules.
* [[Activation energy]], which is defined as the amount of energy required to make the reaction start and carry on spontaneously. Higher activation energy implies that the reactants need more energy to start than a reaction with lower activation energy.
* [[Temperature]], which hastens reactions if raised, since higher temperature increases the energy of the molecules, creating more collisions per unit of time,
* The presence or absence of a [[catalyst]]. Catalysts are substances that make weak bonds with reactants or intermediates and change the pathway (mechanism) of a reaction which in turn increases the speed of a reaction by lowering the [[activation energy]] needed for the reaction to take place. A catalyst is not destroyed or changed during a reaction, so it can be used again.
* For some reactions, the presence of [[electromagnetic radiation]], most notably [[ultraviolet light]], is needed to promote the breaking of bonds to start the reaction.  This is particularly true for reactions involving [[radical (chemistry)|radicals]].

Several theories allow calculating the reaction rates at the molecular level. This field is referred to as reaction dynamics. The rate ''v'' of a [[Rate equation#First-order reactions|first-order reaction]], which could be the disintegration of a substance A, is given by:

:<math chem> v= -\frac {d[\ce{A}]}{dt}= k \cdot [\ce{A}]. </math>

Its integration yields:

:<math chem>\ce{[A]}(t) = \ce{[A]}_{0} \cdot e^{-k\cdot t}. </math>

Here ''k'' is the first-order rate constant, having dimension 1/time, [A](''t'') is the concentration at a time ''t'' and [A]<sub>0</sub> is the initial concentration. The rate of a first-order reaction depends only on the concentration and the properties of the involved substance, and the reaction itself can be described with a characteristic [[half-life]]. More than one time constant is needed when describing reactions of higher order. The temperature dependence of the rate constant usually follows the [[Arrhenius equation]]:

:<math>k = k_0 e^{{-E_a}/{k_{B}T}}</math>

where ''E''<sub>a</sub> is the activation energy and ''k''<sub>B</sub> is the [[Boltzmann constant]]. One of the simplest models of reaction rate is the [[collision theory]]. More realistic models are tailored to a specific problem and include the [[transition state theory]], the calculation of the [[potential energy surface]], the [[Marcus theory]] and the [[RRKM theory|Rice–Ramsperger–Kassel–Marcus (RRKM) theory]].<ref>[[#Atkins|Atkins]], p. 963</ref>

==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 [[Redox#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]].

==Catalysis==
{{main|Catalysis}}
{{Further|Reaction Progress Kinetic Analysis}}
[[File:Activation energy.svg|thumb|right|Schematic potential energy diagram showing the effect of a catalyst in an endothermic chemical reaction. The presence of a catalyst opens a different reaction pathway (in red) with lower activation energy. The final result and the overall thermodynamics are the same.]]
[[File:Pot catalytique vue de la structure.jpg|thumb|Solid heterogeneous catalysts are plated on meshes in ceramic [[catalytic converter]]s in order to maximize their surface area. This exhaust converter is from a [[Peugeot 106]] S2 1100]]

In [[catalysis]], the reaction does not proceed directly, but through a reaction with a third substance known as [[catalyst]]. Although the catalyst takes part in the reaction, forming weak bonds with reactants or intermediates, it is returned to its original state by the end of the reaction and so is not consumed. However, it can be inhibited, deactivated or destroyed by secondary processes. Catalysts can be used in a different phase ([[heterogeneous catalysis|heterogeneous]]) or in the same phase ([[homogeneous catalysis|homogeneous]]) as the reactants. In heterogeneous catalysis, typical secondary processes include [[coking]] where the catalyst becomes covered by [[polymer]]ic side products. Additionally, heterogeneous catalysts can dissolve into the solution in a solid-liquid system or evaporate in a solid–gas system. Catalysts can only speed up the reaction – chemicals that slow down the reaction are called inhibitors.<ref>{{GoldBookRef|title=catalyst|file=C00876}}</ref><ref>{{GoldBookRef|title=inhibitor|file=I03035}}</ref> Substances that increase the activity of catalysts are called promoters, and substances that deactivate catalysts are called catalytic poisons. With a catalyst, a reaction that is kinetically inhibited by high activation energy can take place in the circumvention of this activation energy.

Heterogeneous catalysts are usually solids, powdered in order to maximize their surface area. Of particular importance in heterogeneous catalysis are the [[platinum group]] metals and other transition metals, which are used in [[hydrogenation]]s, [[catalytic reforming]] and in the synthesis of commodity chemicals such as [[nitric acid]] and [[ammonia]]. Acids are an example of a homogeneous catalyst, they increase the nucleophilicity of [[carbonyl]]s, allowing a reaction that would not otherwise proceed with electrophiles. The advantage of homogeneous catalysts is the ease of mixing them with the reactants, but they may also be difficult to separate from the products. Therefore, heterogeneous catalysts are preferred in many industrial processes.<ref>{{cite book | author = Elschenbroich, Christoph | title = Organometallchemie | edition = 6th | publisher = [[Vieweg+Teubner Verlag]] | location = Wiesbaden | year = 2008 | isbn = 978-3-8351-0167-8 | page = 263}}</ref>

==Reactions in organic chemistry==
{{Main|Organic reaction}}
In [[organic chemistry]], in addition to oxidation, reduction or acid-base reactions, a number of other reactions can take place which involves [[covalent bond]]s between carbon atoms or carbon and [[heteroatom]]s (such as oxygen, nitrogen, [[halogen]]s, etc.). Many specific reactions in organic chemistry are [[name reaction]]s designated after their discoverers.

One of the most industrially important reactions is the [[Cracking (chemistry)|cracking]] of heavy [[hydrocarbon]]s at [[Oil refinery|oil refineries]] to create smaller, simpler molecules. This process is used to manufacture [[gasoline]]. Specific types of organic reactions may be grouped by their reaction mechanisms (particularly substitution, addition and elimination) or by the types of products they produce (for example, [[methylation]], [[Polymerization|polymerisation]] and [[halogenation]]).

===Substitution===
In a [[substitution reaction]], a [[functional group]] in a particular [[chemical compound]] is replaced by another group.<ref name=jerry>{{JerryMarch}}</ref> These reactions can be distinguished by the type of substituting species into a [[nucleophilic substitution|nucleophilic]], [[electrophilic substitution|electrophilic]] or [[radical substitution]].
{{multiple image | direction = vertical | image1 = SN1 reaction mechanism.png|width1 = 300|image2 = SN2 reaction mechanism.png|width2 = 300| caption1 = S<sub>N</sub>1 mechanism| caption2 = S<sub>N</sub>2 mechanism}}

In the first type, a [[nucleophile]], an atom or molecule with an excess of electrons and thus a negative charge or [[partial charge]], replaces another atom or part of the "substrate" molecule. The electron pair from the nucleophile attacks the substrate forming a new bond, while the [[leaving group]] departs with an electron pair. The nucleophile may be electrically neutral or negatively charged, whereas the substrate is typically neutral or positively charged. Examples of nucleophiles are [[hydroxide]] ion, [[alkoxide]]s, [[amine]]s and [[halide]]s. This type of reaction is found mainly in [[aliphatic hydrocarbon]]s, and rarely in [[aromatic hydrocarbon]]. The latter have high electron density and enter [[nucleophilic aromatic substitution]] only with very strong [[Polar effect|electron withdrawing groups]]. Nucleophilic substitution can take place by two different mechanisms, [[SN1 reaction|S<sub>N</sub>1]] and [[SN2 reaction|S<sub>N</sub>2]]. In their names, S stands for substitution, N for nucleophilic, and the number represents the [[order (chemistry)|kinetic order]] of the reaction, unimolecular or bimolecular.<ref>{{cite book | author = Hartshorn, S.R. | url = https://books.google.com/books?id=bAo4AAAAIAAJ | title = Aliphatic Nucleophilic Substitution | publisher = [[Cambridge University Press]] | location = London | year = 1973 | isbn = 978-0-521-09801-4 | page = 1}}</ref>
{{multiple image | direction = vertical
| align = right
| width = 120
| image1= Walden-inversion-3D-balls.png
|caption1=The three steps of an [[SN2 reaction|S<sub>N</sub>2 reaction]]. The nucleophile is green and the leaving group is red
|image2=SN2-Walden-before-and-after-horizontal-3D-balls.png
|caption2=S<sub>N</sub>2 reaction causes stereo inversion (Walden inversion)
}}

The S<sub>N</sub>1 reaction proceeds in two steps. First, the [[leaving group]] is eliminated creating a [[carbocation]]. This is followed by a rapid reaction with the nucleophile.<ref>{{Cite journal|author =  Bateman, Leslie C. | author2 = Church, Mervyn G. | author3 = Hughes, Edward D. | author4 = Ingold, Christopher K. | author5 = Taher, Nazeer Ahmed |doi = 10.1039/JR9400000979|title =  188. Mechanism of substitution at a saturated carbon atom. Part XXIII. A kinetic demonstration of the unimolecular solvolysis of alkyl halides. (Section E) a general discussion|year =  1940|journal =  Journal of the Chemical Society|page =  979}}</ref>

In the S<sub>N</sub>2 mechanisms, the nucleophile forms a transition state with the attacked molecule, and only then the leaving group is cleaved. These two mechanisms differ in the [[stereochemistry]] of the products. S<sub>N</sub>1 leads to the non-stereospecific addition and does not result in a chiral center, but rather in a set of [[Cis–trans isomerism|geometric isomers]] (''cis/trans''). In contrast, a reversal ([[Walden inversion]]) of the previously existing stereochemistry is observed in the S<sub>N</sub>2 mechanism.<ref>[[#Bruckner|Brückner]], pp. 63–77</ref>

[[Electrophilic substitution]] is the counterpart of the nucleophilic substitution in that the attacking atom or molecule, an [[electrophile]], has low electron density and thus a positive charge. Typical electrophiles are the carbon atom of [[carbonyl group]]s, carbocations or [[sulfur]] or [[nitronium]] cations. This reaction takes place almost exclusively in aromatic hydrocarbons, where it is called [[electrophilic aromatic substitution]]. The electrophile attack results in the so-called σ-complex, a transition state in which the aromatic system is abolished. Then, the leaving group, usually a proton, is split off and the aromaticity is restored. An alternative to aromatic substitution is electrophilic aliphatic substitution. It is similar to the nucleophilic aliphatic substitution and also has two major types, S<sub>E</sub>1 and S<sub>E</sub>2<ref>[[#Bruckner|Brückner]], pp. 203–206</ref>

[[File:Electrophilic aromatic substitution.svg|center|thumb|648px|Mechanism of electrophilic aromatic substitution]]
{{clear}}
In the third type of substitution reaction, radical substitution, the attacking particle is a [[Radical (chemistry)|radical]].<ref name=jerry/> This process usually takes the form of a [[chain reaction]], for example in the reaction of alkanes with halogens. In the first step, light or heat disintegrates the halogen-containing molecules producing radicals. Then the reaction proceeds as an avalanche until two radicals meet and recombine.<ref>[[#Bruckner|Brückner]], p. 16</ref>
:;<chem>X. + R-H -> X-H + R.</chem>
:;<chem>R. + X2 -> R-X + X.</chem>
::<small> Reactions during the chain reaction of radical substitution </small>

===Addition and elimination===
The [[Addition reaction|addition]] and its counterpart, the [[elimination reaction|elimination]], are reactions that change the number of substituents on the carbon atom, and form or cleave [[covalent bond|multiple bonds]]. [[Double bond|Double]] and [[triple bond]]s can be produced by eliminating a suitable leaving group. Similar to the nucleophilic substitution, there are several possible reaction mechanisms that are named after the respective reaction order. In the E1 mechanism, the leaving group is ejected first, forming a carbocation. The next step, the formation of the double bond, takes place with the elimination of a proton ([[deprotonation]]). The leaving order is reversed in the E1cb mechanism, that is the proton is split off first. This mechanism requires the participation of a base.<ref>[[#Bruckner|Brückner]], p. 192</ref> Because of the similar conditions, both reactions in the E1 or E1cb elimination always compete with the S<sub>N</sub>1 substitution.<ref>[[#Bruckner|Brückner]], p. 183</ref>

{{multiple image
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| image1 = E1-mechanism.svg
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| caption1 = E1 elimination
| image2 = E1cb-mechanism.svg
| width2 = 400
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| caption2 = E1cb elimination
| footer =
}}
{{clear}}
[[File:E2-mechanism.svg|thumb|300px|E2 elimination]]

The E2 mechanism also requires a base, but there the attack of the base and the elimination of the leaving group proceed simultaneously and produce no ionic intermediate. In contrast to the E1 eliminations, different stereochemical configurations are possible for the reaction product in the E2 mechanism, because the attack of the base preferentially occurs in the anti-position with respect to the leaving group. Because of the similar conditions and reagents, the E2 elimination is always in competition with the S<sub>N</sub>2-substitution.<ref>[[#Bruckner|Brückner]], p. 172</ref>

[[File:HBr-addition.svg|thumb|300px|Electrophilic addition of hydrogen bromide]]
The counterpart of elimination is an addition where double or triple bonds are converted into single bonds. Similar to substitution reactions, there are several types of additions distinguished by the type of the attacking particle. For example, in the [[electrophilic addition]] of [[hydrogen bromide]], an electrophile (proton) attacks the double bond forming a [[carbocation]], which then reacts with the nucleophile (bromine). The carbocation can be formed on either side of the double bond depending on the groups attached to its ends, and the preferred configuration can be predicted with the [[Markovnikov's rule]].<ref>[[#Wiberg|Wiberg]], pp. 950, 1602</ref> This rule states that "In the heterolytic addition of a polar molecule to an alkene or alkyne, the more electronegative (nucleophilic) atom (or part) of the polar molecule becomes attached to the carbon atom bearing the smaller number of hydrogen atoms."<ref>{{GoldBookRef|title=Markownikoff rule|file=M03707}}</ref>

If the addition of a functional group takes place at the less substituted carbon atom of the double bond, then the electrophilic substitution with acids is not possible. In this case, one has to use the [[hydroboration–oxidation reaction]], wherein the first step, the [[boron]] atom acts as electrophile and adds to the less substituted carbon atom. In the second step, the nucleophilic [[hydroperoxide]] or halogen [[anion]] attacks the boron atom.<ref>[[#Bruckner|Brückner]], p. 125</ref>

While the addition to the electron-rich alkenes and alkynes is mainly electrophilic, the [[nucleophilic addition]] plays an important role in the carbon-heteroatom multiple bonds, and especially its most important representative, the carbonyl group. This process is often associated with elimination so that after the reaction the carbonyl group is present again. It is, therefore, called an addition-elimination reaction and may occur in carboxylic acid derivatives such as chlorides, esters or anhydrides. This reaction is often catalyzed by acids or bases, where the acids increase the electrophilicity of the carbonyl group by binding to the oxygen atom, whereas the bases enhance the nucleophilicity of the attacking nucleophile.<ref>{{cite book | author = Latscha, Hans Peter | author2 = Kazmaier, Uli | author3 = Klein, Helmut Alfons | title = Organische Chemie: Chemie-basiswissen II | volume = 2 | edition = 6th | language = de | publisher = [[Springer Science+Business Media|Springer]] | year = 2008 | isbn = 978-3-540-77106-7 | page = 273}}</ref>

[[File:H-Add-El.Mechanismus.PNG|thumb|center|500px|Acid-catalyzed addition-elimination mechanism]]
{{clear}}

[[Nucleophilic addition]] of a [[carbanion]] or another [[nucleophile]] to the double bond of an [[Α,β-unsaturated carbonyl compound|alpha, beta-unsaturated carbonyl compound]] can proceed via the [[Michael reaction]], which belongs to the larger class of [[conjugate addition]]s. This is one of the most useful methods for the mild formation of C–C bonds.<ref>{{cite book|year=2004|doi=10.1002/0471264180|title=Organic Reactions|isbn=978-0-471-26418-7 |editor-last1=Denmark |editor-first1=Scott E. }}</ref><ref>{{cite web|title = Chapter 18: Enols and Enolates — The Michael Addition reaction|author = Hunt, Ian |publisher = University of Calgary|url = http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch18/ch18-4-3.html}}</ref><ref>[[#Bruckner|Brückner]], p. 580</ref>

Some additions which can not be executed with nucleophiles and electrophiles can be succeeded with free radicals. As with the free-radical substitution, the [[radical addition]] proceeds as a chain reaction, and such reactions are the basis of the [[Radical polymerization|free-radical polymerization]].<ref>{{cite book | author = Lechner, Manfred | author2 = Gehrke, Klaus | author3 = Nordmeier, Eckhard | title = Macromolecular Chemistry | edition = 3rd | publisher = [[Birkhäuser]] | location =  Basel | year = 2003 | isbn = 978-3-7643-6952-1 | pages = 53–65}}</ref>

===Other organic reaction mechanisms===
[[File:Cope Rearrangement Scheme.png|left|thumb|The Cope rearrangement of 3-methyl-1,5-hexadiene]]
{{multiple image | direction = vertical
| align = right
| width = 220
| image1= Diels Alder Mechanismus.svg
|caption1=Mechanism of a Diels-Alder reaction
| image2= Diels Alder Orbitale.svg
|caption2=Orbital overlap in a Diels-Alder reaction}}

In a [[rearrangement reaction]], the carbon skeleton of a [[molecule]] is rearranged to give a [[structural isomer]] of the original molecule. These include [[Sigmatropic reaction|hydride shift]] reactions such as the [[Wagner-Meerwein rearrangement]], where a [[hydrogen]], [[alkyl]] or [[aryl]] group migrates from one carbon to a neighboring carbon. Most rearrangements are associated with the breaking and formation of new carbon-carbon bonds. Other examples are [[sigmatropic reaction]] such as the [[Cope rearrangement]].<ref>{{cite book | author = Fox, Marye Anne | author2 = Whitesell, James K. | url = https://books.google.com/books?id=xx_uIP5LgO8C&pg=PA699 | title = Organic chemistry | edition = Third | publisher = [[Jones & Bartlett Learning|Jones & Bartlett]] | year = 2004 | isbn = 978-0-7637-2197-8 | page = 699}}</ref>

Cyclic rearrangements include [[cycloaddition]]s and, more generally, [[pericyclic reaction]]s, wherein two or more double bond-containing molecules form a cyclic molecule. An important example of cycloaddition reaction is the [[Diels–Alder reaction]] (the so-called [4+2] cycloaddition) between a conjugated [[diene]] and a substituted [[alkene]] to form a substituted [[cyclohexene]] system.<ref>{{Cite journal| volume = 460| pages = 98–122| year = 1928 | doi = 10.1002/jlac.19284600106| journal = Justus Liebig's Annalen der Chemie| title = Synthesen in der hydroaromatischen Reihe| first2 = K.| last2 = Alder| last1 = Diels| first1 = O.}}</ref>

Whether a certain cycloaddition would proceed depends on the electronic orbitals of the participating species, as only orbitals with the same sign of [[wave function]] will overlap and interact constructively to form new bonds. Cycloaddition is usually assisted by light or heat. These perturbations result in a different arrangement of electrons in the excited state of the involved molecules and therefore in different effects. For example, the [4+2] Diels-Alder reactions can be assisted by heat whereas the [2+2] cycloaddition is selectively induced by light.<ref>[[#Bruckner|Brückner]], pp. 637–647</ref> Because of the orbital character, the potential for developing [[stereochemistry|stereoisomeric]] products upon cycloaddition is limited, as described by the [[Woodward–Hoffmann rules]].<ref>{{Cite journal | doi = 10.1021/ja01080a054| title = Stereochemistry of Electrocyclic Reactions| journal = Journal of the American Chemical Society| volume = 87| issue = 2| pages = 395–397| year = 1965| last1 = Woodward | first1 = R.B.| last2 = Hoffmann | first2 = R. | bibcode = 1965JAChS..87..395W}}</ref>

==Biochemical reactions==
[[File:Induced fit diagram.svg|thumb|left|380px|Illustration of the induced fit model of enzyme activity]]
[[Biochemistry|Biochemical reactions]] are mainly controlled by complex [[protein]]s called [[enzymes]], which are usually specialized to [[Catalysis|catalyze]] only a single, specific reaction. The reaction takes place in the [[active site]], a small part of the enzyme which is usually found in a cleft or pocket lined by [[amino acid]] residues, and the rest of the enzyme is used mainly for stabilization. The catalytic action of enzymes relies on several mechanisms including the molecular shape ("induced fit"), bond strain, proximity and orientation of molecules relative to the enzyme, proton donation or withdrawal (acid/base catalysis), electrostatic interactions and many others.<ref>{{cite book | author = Karlson, Peter | author2 = Doenecke, Detlef | author3 = Koolman, Jan | author4 = Fuchs, Georg | author5 = Gerok, Wolfgang | title = Karlson Biochemistry and Pathobiochemistry | edition = 16th | url = https://books.google.com/books?id=8Eu_Hy8aVzMC&pg=PP1| language = de | publisher = [[Thieme Medical Publishers|Thieme]] | year = 2005 | isbn = 978-3-13-357815-8 | pages = 55–56}}</ref>

The biochemical reactions that occur in living organisms are collectively known as [[metabolism]]. Among the most important of its mechanisms is the [[anabolism]], in which different [[DNA]] and enzyme-controlled processes result in the production of large molecules such as [[protein]]s and [[carbohydrates]] from smaller units.<ref>{{GoldBookRef|title=anabolism|file=A00314}}</ref> [[Bioenergetics]] studies the sources of energy for such reactions. Important energy sources are [[glucose]] and [[Dioxygen in biological reactions|oxygen]], which can be produced by plants via [[photosynthesis]] or assimilated from food and air, respectively. All organisms use this energy to produce [[adenosine triphosphate]] (ATP), which can then be used to energize other reactions. Decomposition of organic material by [[Fungus|fungi]], [[bacteria]] and other [[Microorganism|micro-organisms]] is also within the scope of [[biochemistry]].

==Applications==
[[File:Velp-thermitewelding-1.jpg|thumb|Thermite reaction proceeding in railway welding. Shortly after this, the liquid iron flows into the mould around the rail gap.]]
Chemical reactions are central to [[chemical engineering]], where they are used for the synthesis of new compounds from natural raw materials such as [[petroleum]], mineral [[ore]]s, and [[oxygen]] in air. It is essential to make the reaction as efficient as possible, maximizing the yield and minimizing the number of reagents, energy inputs and waste. [[Catalyst]]s are especially helpful for reducing the energy required for the reaction and increasing its [[reaction rate]].<ref>{{cite book | author = Emig, Gerhard | author2 =  Klemm, Elias | title = Technical Chemistry | edition = 5th | language = de | publisher = [[Springer Science+Business Media|Springer]] | year = 2005 | isbn = 978-3-540-23452-4 | pages = 33–34}}</ref><ref>{{Cite journal | doi = 10.1126/science.1962206| title = The atom economy – a search for synthetic efficiency| journal = Science| volume = 254| issue = 5037| pages = 1471–1477| year = 1991| last1 = Trost | first1 = B. | pmid=1962206| bibcode = 1991Sci...254.1471T}}</ref>

Some specific reactions have their niche applications. For example, the [[thermite]] reaction is used to generate light and heat in [[pyrotechnics]] and [[welding]]. Although it is less controllable than the more conventional [[Oxy-fuel welding and cutting|oxy-fuel welding]], [[arc welding]] and [[flash welding]], it requires much less equipment and is still used to mend rails, especially in remote areas.<ref>{{cite encyclopedia | editor = John J. McKetta | author = Weismantel, Guy E | year = 1999 | url = https://books.google.com/books?id=MfjDlUe8Kc0C&pg=PA109 | title= Encyclopedia of Chemical Processing and Design | volume = 67 | publisher = [[CRC Press]] | page = 109 | isbn = 978-0-8247-2618-8}}</ref>

==Monitoring==
Mechanisms of monitoring chemical reactions depend strongly on the reaction rate. Relatively slow processes can be analyzed in situ for the concentrations and identities of the individual ingredients. Important tools of real-time analysis are the measurement of [[pH]] and analysis of optical absorption (color) and emission spectra. A less accessible but rather efficient method is the introduction of a radioactive isotope into the reaction and monitoring how it changes over time and where it moves to; this method is often used to analyze the redistribution of substances in the human body. Faster reactions are usually studied with [[ultrafast laser spectroscopy]] where utilization of [[Femtochemistry|femtosecond]] [[laser]]s allows short-lived transition states to be monitored at a time scaled down to a few femtoseconds.<ref>[[#Atkins|Atkins]], p. 987</ref>

==See also==
{{Portal|Chemistry}}
{{wikiquote}}
* [[Chemical equation]]
* Chemical reaction
** [[Substrate (chemistry)|Substrate]]
** [[Reagent]]
** [[Catalyst]]
** [[Product (chemistry)|Product]]
* [[Chemical reaction model]]
* [[Chemist]]
* [[Chemistry]]
* [[Combustion]]
* [[Limiting reagent]]
* [[List of organic reactions]]
* [[Mass balance]]
* [[Microscopic reversibility]]
* [[Organic reaction]]
* [[Reaction progress kinetic analysis]]
* [[Reversible reaction]]

==References==
{{reflist}}

==Bibliography==
* {{cite book|ref=Atkins|author1=Atkins, Peter W.  |author2=Julio de Paula|title=Physical Chemistry|edition = 4th|publisher=[[Wiley-VCH]]|location=Weinheim|year=2006|isbn= 978-3-527-31546-8}}
* {{cite book|ref=Brock|author=Brock, William H.|url=https://books.google.com/books?id=AJ-c8py7t6gC&pg=PA459 |title=Viewegs Geschichte der Chemie|language = de|publisher=[[Vieweg+Teubner Verlag|Vieweg]]|location=Braunschweig |year=1997|isbn=978-3-540-67033-9}}
* {{cite book|ref=Bruckner|author=Brückner, Reinhard| title = Reaktionsmechanismen |edition = 3rd|language = de|publisher=Spektrum Akademischer Verlag|location=München|year= 2004|isbn=978-3-8274-1579-0}}
* {{cite book|ref=Wiberg|author=Wiberg, Egon, Wiberg, Nils and Holleman, Arnold Frederick |url=https://books.google.com/books?id=Mtth5g59dEIC&pg=PA287 |title=Inorganic chemistry|publisher= [[Academic Press]]|year= 2001 |isbn= 978-0-12-352651-9}}
* {{cite EB1911|wstitle=Chemical Action|volume=6|pages=26–33|short=1}}

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{{Chemical synthesis}}
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{{Organic reactions}}
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