## Synthesis Reactions

Synthesis reactions always yield one product. Reversing a synthesis reaction will give you a decomposition reaction.
The general form of a synthesis reaction is A + B → AB. Synthesis reactions "put things together".
 ${\displaystyle 2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}}$ This is the most well-known example of a synthesis reaction—the formation of water via the fusion of hydrogen gas and oxygen gas. ${\displaystyle 2{\hbox{Na}}_{(s)}+{\hbox{Cl}}_{2(g)}\to 2{\hbox{NaCl}}_{(s)}}$ Another example of a synthesis reaction is the formation of sodium chloride (table salt).
Because of the very high reactivities of sodium metal and chlorine gas, this reaction releases a tremendous amount of heat and light energy. Recall that atoms release energy as they become stable, and consider the octet rule when determining why this reaction is so favorable.

## Decomposition Reactions

These are the opposite of synthesis reactions, with the format AB → A + B. Decomposition reactions "take things apart". Just as synthesis reactions can only form one product, decomposition reactions can only start with one reactant. Compounds that are unstable will decompose quickly without outside assistance.
 ${\displaystyle 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}{\xrightarrow {electricity}}2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}}$ One example is the electrolysis of water (passing water through electrical current) to form hydrogen gas and oxygen gas. ${\displaystyle 2{\hbox{H}}_{2}{\hbox{O}}_{2(l)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{O}}_{2(g)}}$ Hydrogen peroxide slowly decomposes into water and oxygen because it is somewhat unstable. The process is sped up by the energy from light, so hydrogen peroxide is stored in dark containers to slow down the decomposition. ${\displaystyle {\hbox{H}}_{2}{\hbox{CO}}_{3(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{C}}{\hbox{O}}_{2(g)}}$ Carbonic acid is the carbonation that is dissolved in soda. It decomposes into carbon dioxide and water, which is why an opened drink will not lose its fizz.

## Single Replacement Reactions

Single replacement reactions, also called single displacement, swap one component with another, in the format AB + C → AC + B.
Adding hydrochloric acid to zinc will cause a gas to bubble out:
${\displaystyle {\hbox{Zn}}_{(s)}+2{\hbox{HCl}}_{(aq)}\to {\hbox{ZnCl}}_{2(aq)}+{\hbox{H}}_{2(g)}}$

## Double Replacement Reactions

For further information, see "Double Displacement Reaction"
In these reactions, also known as "double displacement reactions", two compounds swap components, in the format AB + CD → AD + CB

## Double Displacement Reaction

This is also called an "exchange". Here are the examples below:
1.) HCl + NaOH ----> NaCl + H2O

### Precipitation

A precipitation reaction occurs when an ionic substance comes out of solution and forms an insoluble (or slightly soluble) solid. The solid which comes out of solution is called a precipitate. This can occur when two soluble salts (ionic compounds) are mixed and form an insoluble one—the precipitate.
 ${\displaystyle {\hbox{2Pb}}({\hbox{NO}}_{3})_{2(aq)}+heat_{(aq)}\to {\hbox{2PbO}}_{(s)}+4{\hbox{NO}}_{2(aq)}+{O}_{2}}$ An example is lead nitrate mixed with potassium iodide, which forms a bright yellow precipitate of lead iodide. ${\displaystyle {\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{NO}}_{3(aq)}^{-}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{NO}}_{3(aq)}^{-}}$ Note that the lead iodide is formed as a solid. The previous equation is written in molecular form, which is not the best way of describing the reaction. Each of the elements really exist in solution as individual ions, not bonded to each other (as in potassium iodide crystals). If we write the above as an ionic equation, we get a much better idea of what is actually happening. ${\displaystyle {\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}}$ Notice the like terms on both sides of the equation. These are called spectator ionsbecause they do not participate in the reaction. They can be ignored, and the net ionic equation is written.
In the solution, there exists both lead and iodide ions. Because lead iodide is insoluble, they spontaneously crystallise and form the precipitate.

### Acid-Base Neutralization

In simple terms, an acid is a substance which can lose a H+ ion (i.e. a proton) and a base is a substance which can accept a proton. When equal amounts of an acid and base react, they neutralize each other, forming species which aren't as acidic or basic.
 ${\displaystyle {\hbox{HCl}}_{(aq)}+{\hbox{NaOH}}_{(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{NaCl}}_{(aq)}}$ For example, when hydrochloric acid and sodium hydroxide react, they form water and sodium chloride (table salt). ${\displaystyle {\hbox{H}}_{(aq)}^{+}+{\hbox{OH}}_{(aq)}^{-}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}}$ Again, we get a clearer picture of what's happening if we write a net ionic equation.
Acid base reactions often happen in aqueous solution, but they can also occur in the gaseous state. Acids and bases will be discussed in much greater detail in the acids and bases section.

## Combustion

The combustion of methane (releasing heat and light)
Combustion, better known as burning, is the combination of a substance with oxygen. The products are carbon dioxide, water, and possible other waste products. Combustion reactions release large amounts of heat. C3H8, better known as propane, undergoes combustion. The balanced equation is:
${\displaystyle {\hbox{C}}_{3}{\hbox{H}}_{8}+5{\hbox{O}}_{2}\to 3{\hbox{CO}}_{2}+4{\hbox{H}}_{2}{\hbox{O}}}$
Combustion is similar to a decomposition reaction, except that oxygen and heat are required for it to occur. If there is not enough oxygen, the reaction may not occur. Sometimes, with limited oxygen, the reaction will occur, but it produces carbon monoxide (CO) or even soot. In that case, it is called incomplete combustion. If the substances being burned contain atoms other than hydrogen and oxygen, then waste products will also form. Coal is burned for heating and energy purposes, and it contains sulfur. As a result, sulfur dioxide is released, which is a pollutant. Coal with lower sulfur content is more desirable, but more expensive, because it will release less of the sulfur-based pollutants.

## Organic Reactions

This is carboxylic acid. All functional groups end with an "R"—a placeholder for the rest of the molecule.
Organic reactions occur between organic molecules (molecules containing carbon and hydrogen). Since there is a virtually unlimited number of organic molecules, the scope of organic reactions is very large. However, many of the characteristics of organic molecules are determined byfunctional groups—small groups of atoms that react in predictable ways.
Another key concept in organic reactions is Lewis basicity. Parts of organic molecules can be electrophillic (electron-loving) or nucleophillic (nucleus, or positive loving). Nucleophillic regions have an excess of electrons—they act as Lewis bases—whereas electrophillic areas are electron deficient and act as Lewis acids. The nucleophillic and electrophillic regions attract and react with each other (needless to say, this has inspired many terrible organic chemistry jokes).
Organic reactions are beyond the scope of this book, and are covered in more detail in Organic Chemistry. However, most organic substances can undergo replacement reactions and combustion reactions, as you have already learned.

## Redox

The formation of hydrogen fluoride causes fluoride to oxidize and hydrogen to reduce.
Redox is an abbreviation of reduction/oxidation reactions. This is exactly what happens in a redox reaction, one species is reducedand another is oxidized. Reduction involves a gain of electrons and oxidation involves a loss, so a redox reaction is one in which electrons are transferred between species. Reactions where something is "burnt" (burning means being oxidised) are examples of redox reactions, however, oxidation reactions also occur in solution, which is very useful and forms the basis of electrochemistry.
Redox reactions are often written as two half-reactions showing the reduction and oxidation processes separately. These half-reactions are balanced (by multiplying each by a coefficient) and added together to form the full equation. When magnesium is burnt in oxygen, it loses electrons (it is oxidised). Conversely, the oxygen gains electrons from the magnesium (it is reduced).
${\displaystyle {\begin{matrix}{\hbox{Mg}}&\to &{\hbox{Mg}}^{2+}+2e^{-}&\times 2\\{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{O}}^{2-}&\times 1\\2{\hbox{Mg}}+{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{MgO}}+4e^{-}&\ \\\end{matrix}}}$
Redox reactions will be discussed in greater detail in the redox section.

## Equations

As seen from the equation CH
4
+ 2 O
2
→ CO
2
+ 2 H
2
O
, a coefficient of 2 must be placed before the oxygen gas on the reactants side and before the 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 article: Chemical equation
Chemical equations are used to graphically illustrate chemical reactions. They consist of chemicalor structural formulas 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".[7] The tip of the arrow points in the direction in which the reaction proceeds. A double arrow (⇌) pointing in opposite directions is used for 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 (${\displaystyle {\ce {A,B,C}}}$ and ${\displaystyle {\ce {D}}}$ in a schematic example below) by the appropriate integers a, b, c and d.[8]
${\displaystyle {\ce {{{\mathit {a}}A}+{\mathit {b}}B->{{\mathit {c}}C}+{\mathit {d}}D}}}$
More elaborate reactions are represented by reaction schemes, which in addition to starting materials and products show important intermediates or transition states. 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.
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.[9]

## Elementary reactions

The elementary reaction is the smallest division into which a chemical reaction can be decomposed, it has no intermediate products.[10] 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.[11]

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 an isomerization or a 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 isomerization, in which the cis-form of a compound converts to the trans-form or vice versa.[12]
In a typical dissociation reaction, a bond in a molecule splits (ruptures) resulting in two molecular fragments. The splitting can be homolytic or heterolytic. In the first case, the bond is divided so that each product retains an electron and becomes a neutral radical. In the second case, both electrons of the chemical bond remain with one of the products, resulting in charged ions. Dissociation plays an important role in triggering chain reactions, such as hydrogen–oxygen orpolymerization reactions.
${\displaystyle {\ce {AB->{A}+{B}}}}$
Dissociation of a molecule AB into fragments A and B
For bimolecular reactions, two molecules collide and react with each other. Their merger is called chemical synthesis or an addition reaction.
${\displaystyle {\ce {{A}+{B}->AB}}}$
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 metathesis.
${\displaystyle {\ce {{HA}+{B}->{A}+{HB}}}}$
for example
${\displaystyle {\ce {{NaCl}+{AgNO3}->{NaNO3}+{AgCl(v)}}}}$

## Chemical equilibrium

Main article: 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 reaction rates. These rates depend on the concentration and therefore change with 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 such parameters as temperature, pressure and the materials involved, and is determined by the minimum free energy. In equilibrium, the Gibbs free energy 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 the fewer moles of gas.[13]
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 energy. The associated free energy of the reaction is composed of two different thermodynamic quantities, enthalpy and entropy:[14]
${\displaystyle \Delta G=\Delta H-T\cdot \Delta S}$.
G: free energy, H: enthalpy, T: temperature, S: entropy, Δ: difference(change between original and product)
Reactions can be exothermic, where ΔH is negative and energy is released. Typical examples of exothermic reactions are 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 reaction products, which have high entropy. Since the entropy increases with temperature, many endothermic reactions preferably take place at high temperatures. On the contrary, many exothermic reactions such as crystallization occur at low temperatures. Changes in temperature can sometimes reverse the sign of the enthalpy of a reaction, as for the carbon monoxide reduction of molybdenum dioxide:
${\displaystyle {\ce {{2CO(g)}+{MoO2(s)}->{2CO2(g)}+{Mo(s)}}}}$${\displaystyle \Delta H^{o}=+21.86\ {\text{kJ at 298 K}}}$
This reaction to form carbon dioxide and molybdenum is endothermic at low temperatures, becoming less so with increasing temperature.[15] ΔH° is zero at 1855 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
${\displaystyle {\ce {{CO(g)}+{H2O({v})}<=>{CO2(g)}+{H2(g)}}}}$
is favored by low temperatures, but its reverse is favored by high temperature. The shift in reaction direction tendency occurs at 1100 K.[15]
Reactions can also be characterized by the internal energy which takes into account changes in the entropy, volume and chemical potential. The latter depends, among other things, on the activities of the involved substances.[16]
${\displaystyle {d}U=T\cdot {d}S-p\cdot {d}V+\mu \cdot {d}n}$
U: internal energy, S: entropy, p: pressure, μ: chemical potential, n: number of molecules, dsmall change sign

## Kinetics

The speed at which reactions takes place is studied by 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 time. Some reactions, however, have rates that are independent of reactant concentrations. These are called 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 a lower activation energy.
• Temperature, which hastens reactions if raised, since higher temperature increases the energy of the molecules, creating more collisions per unit time,
• The presence or absence of a catalyst. Catalysts are substances which 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 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 first-order reaction, which could be disintegration of a substance A, is given by:
${\displaystyle v=-{\frac {d[{\ce {A}}]}{dt}}=k\cdot [{\ce {A}}].}$
Its integration yields:
${\displaystyle {\ce {[A]}}(t)={\ce {[A]}}_{0}\cdot e^{-k\cdot t}.}$
Here k is first-order rate constant having dimension 1/time, [A](t) is concentration at a time t and [A]0 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 the 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:
${\displaystyle k=k_{0}e^{{-E_{a}}/{k_{B}T}}}$
where Ea is the activation energy and kB 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 Rice–Ramsperger–Kassel–Marcus (RRKM) theory.[17]

## Reaction types

### Four basic types

Representation of four basic chemical reactions types: synthesis, decomposition, single replacement and double replacement.

#### Synthesis

Main article: 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:
${\displaystyle {\ce {{A}+{B}->AB}}}$
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:
${\displaystyle {\ce {{8Fe}+S8->8FeS}}}$
Another example is simple hydrogen gas combined with simple oxygen gas to produce a more complex substance, such as water.[18]

#### Decomposition

Main article: 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[18][19]
${\displaystyle {\ce {AB->{A}+{B}}}}$
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas:
${\displaystyle {\ce {2H2O->{2H2}+{O2}}}}$

#### Single replacement

In a single replacement reaction, a single uncombined element replaces another in a compound; in other words, one element trades places with another element in a compound[18] These reactions come in the general form of:
${\displaystyle {\ce {{A}+{BC}->{AC}+{B}}}}$
One example of a single displacement reaction is when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas:
${\displaystyle {\ce {{Mg}+{2H2O}->{Mg(OH)2}+{H2\uparrow }}}}$

#### Double replacement

In a double replacement reaction, the anions and cations of two compounds switch places and form two entirely different compounds.[18] These reactions are in the general form:[19]
${\displaystyle {\ce {{AB}+{CD}->{AD}+{CB}}}}$
For example, when barium chloride (BaCl2) and magnesium sulfate (MgSO4) react, the SO42− anion switches places with the 2Cl anion, giving the compounds BaSO4 and MgCl2.
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:
${\displaystyle {\ce {{Pb(NO3)2}+2KI->PbI2(v)+2KNO3}}}$

### Oxidation and reduction

Illustration of a redox reaction

Sodium chloride is formed through the redox reaction of sodium metal and chlorine gas
Redox reactions can be understood in terms of 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, 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).[20][21]
In the following redox reaction, hazardous sodium metal reacts with toxic chlorine gas to form the ionic compound sodium chloride, or common table salt:
${\displaystyle {\ce {{2Na(s)}+{Cl2(g)}->2NaCl(s)}}}$
In the reaction, sodium metal goes from an oxidation state of 0 (as it is 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 (it is 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 reduction in the other species and is considered the reducing agent.
Which of the involved reactants would be reducing or oxidizing agent can be predicted from the electronegativity of their elements. Elements with low electronegativity, such as most metals, easily donate electrons and oxidize – they are reducing agents. On the contrary, many ions with high oxidation numbers, such as H
2
O
2
MnO
4
CrO
3
Cr
2
O2−
7
OsO
4
can gain one or two extra electrons and are strong oxidizing agents.
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.[22]
An important class of redox reactions are the electrochemical reactions, where electrons from the power supply are used as the reducing agent. These reactions are particularly important for the production of chemical elements, such as chlorine[23] or aluminium. The reverse process in which electrons are released in redox reactions and can be used as electrical energy is possible and used in batteries.

### Complexation

Ferrocene – an iron atom sandwiched between two C5H5 ligands
In complexation reactions, several ligands react with a metal atom to form a coordination complex. This is achieved by providing lone pairs of the ligand into empty orbitals of the metal atom and forming dipolar bonds. The ligands are Lewis bases, 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 shells of a transition metal will collectively accommodate 18 electrons, 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.[24]

### Acid-base reactions

In the Brønsted–Lowry acid–base theory, an acid-base reaction involves a transfer of protons (H+) from one species (the acid) to another (thebase). When a proton is removed from an acid, the resulting species is termed that acid's conjugate base. When the proton is accepted by a base, the resulting species is termed that base's conjugate acid.[25] In other words, acids act as proton donors and bases act as proton acceptors according to the following equation:
${\displaystyle {\ce {\underbrace {HA} _{acid}+\underbrace {B} _{base}<=>\underbrace {A^{-}} _{conjugated\ base}+\underbrace {HB+} _{conjugated\ acid}}}}$
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 and base dissociation constants (Ka and Kb) of the involved substances. A special case of the acid-base reaction is the neutralization where an acid and a base, taken at exactly same amounts, form a neutral salt.
Acid-base reactions can have different definitions depending on the acid-base concept employed. Some of the most common are:
• Arrhenius definition: Acids dissociate in water releasing H3O+ ions; bases dissociate in water releasing OH ions.
• Brønsted-Lowry definition: Acids are proton (H+) donors, bases are proton acceptors; this includes the Arrhenius definition.
• Lewis definition: Acids are electron-pair acceptors, bases are electron-pair donors; this includes the Brønsted-Lowry definition.

### Precipitation

Precipitation
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[26] and forms an insoluble salt. This process can be assisted by adding a precipitating agent or by removal of the solvent. Rapid precipitation results in an amorphous or microcrystalline residue and slow process can yield single crystals. The latter can also be obtained by recrystallization from microcrystalline salts.[27]

### 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.[28]

### Reactions at the solid|gas interface

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.[29][30] Reactions at the solid|gas interface are in some cases related to catalysis.

### Photochemical reactions

In this Paterno–Büchi reaction, a photoexcited carbonyl group is added to an unexcited olefin, yielding anoxetane.
In photochemical reactions, atoms and molecules absorb energy (photons) of the illumination light and convert into an excited state. They can then release this energy by breaking chemical bonds, thereby producing radicals. Photochemical reactions include hydrogen–oxygen reactions, radical polymerizationchain reactions and rearrangement reactions.[31]
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 vision is initiated by a photochemical reaction of rhodopsin.[12] In fireflies, an enzyme in the abdomen catalyzes a reaction that results in bioluminescence.[32] Many significant photochemical reactions, such as ozone formation, occur in the Earth atmosphere and constitute atmospheric chemistry.

## Catalysis

Main article: Catalysis
Further information: Reaction Progress Kinetic Analysis

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 a lower activation energy. The final result and the overall thermodynamics are the same.

Solid heterogeneous catalysts are plated on meshes in ceramic catalytic converters 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 reaction with a third substance known as catalyst. Although the catalyst takes part in the reaction, 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) or in the same phase (homogeneous) as the reactants. In heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric 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.[33][34] Substances that increase the activity of catalysts are called promoters, and substances that deactivate catalysts are called catalytic poisons. With a catalyst, a reaction which is kinetically inhibited by a high activation energy can take place in 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 hydrogenationscatalytic 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 carbonyls, 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.[35]

## Reactions in organic chemistry

Main article: Organic reaction
In organic chemistry, in addition to oxidation, reduction or acid-base reactions, a number of other reactions can take place which involvecovalent bonds between carbon atoms or carbon and heteroatoms (such as oxygen, nitrogen, halogens, etc.). Many specific reactions in organic chemistry are name reactions designated after their discoverers.

### Substitution

In a substitution reaction, a functional group in a particular chemical compound is replaced by another group.[36] These reactions can be distinguished by the type of substituting species into a nucleophilicelectrophilic or radical substitution.
SN1 mechanism
SN2 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,alkoxidesamines and halides. This type of reaction is found mainly in aliphatic hydrocarbons, and rarely in aromatic hydrocarbon. The latter have high electron density and enter nucleophilic aromatic substitution only with very strongelectron withdrawing groups. Nucleophilic substitution can take place by two different mechanisms, SN1 and SN2. In their names, S stands for substitution, N for nucleophilic, and the number represents the kinetic order of the reaction, unimolecular or bimolecular.[37]
The three steps of anSN2 reaction. The nucleophile is green and the leaving group is red
SN2 reaction causes stereo inversion (Walden inversion)
The SN1 reaction proceeds in two steps. First, the leaving group is eliminated creating a carbocation. This is followed by a rapid reaction with the nucleophile.[38]
In the SN2 mechanism, 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. SN1 leads to the non-stereospecific addition and does not result in a chiral center, but rather in a set of geometric isomers (cis/trans). In contrast, a reversal (Walden inversion) of the previously existing stereochemistry is observed in the SN2 mechanism.[39]
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 groups, 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, SE1 and SE2[40]
Mechanism of electrophilic aromatic substitution
In the third type of substitution reaction, radical substitution, the attacking particle is a radical.[36] 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 the radicals. Then the reaction proceeds as an avalanche until two radicals meet and recombine.[41]
${\displaystyle {\ce {{X.}+{R-H}->{X-H}+{R.}}}}$
${\displaystyle {\ce {{R.}+{X2}->{R{-}X}+{X.}}}}$
Reactions during the chain reaction of radical substitution

The addition and its counterpart, the elimination, are reactions which change the number of substitutents on the carbon atom, and form or cleave multiple bondsDouble andtriple bonds can be produced by eliminating a suitable leaving group. Similar to the nucleophilic substitution, there are several possible reaction mechanisms which are named after the respective reaction order. In the E1 mechanism, the leaving group is ejected first, forming a carbocation. The next step, formation of the double bond, takes place with 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 participation of a base.[42] Because of the similar conditions, both reactions in the E1 or E1cb elimination always compete with the SN1 substitution.[43]
E1 elimination
E1cb elimination

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 SN2-substitution.[44]

The counterpart of elimination is the addition where double or triple bonds are converted into single bonds. Similar to the 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.[45] 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."[46]
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, where in the first step, the boron atom acts as electrophile and adds to the less substituted carbon atom. At the second step, the nucleophilic hydroperoxide or halogen anion attacks the boron atom.[47]
While the addition to the electron-rich alkenes and alkynes is mainly electrophilic, the nucleophilic addition plays an important role for the carbon-heteroatom multiple bonds, and especially its most important representative, the carbonyl group. This process is often associated with an elimination, so that after the reaction the carbonyl group is present again. It is therefore called 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 by the electrophilicity of the carbonyl group by binding to the oxygen atom, whereas the bases enhance the nucleophilicity of the attacking nucleophile.[48]

Nucleophilic addition of a carbanion or another nucleophile to the double bond of an alpha, beta unsaturated carbonyl compound can proceed via the Michael reaction, which belongs to the larger class of conjugate additions. This is one of the most useful methods for the mild formation of C–C bonds.[49][50][51]
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 additionproceeds as a chain reaction, and such reactions are the basis of the free-radical polymerization.[52]

### Other organic reaction mechanisms

Mechanism of a Diels-Alder reaction
Orbital overlap in a Diels-Alder reaction
In a rearrangement reaction, the carbon skeleton of a molecule is rearranged to give astructural isomer of the original molecule. These include hydride shift reactions such as theWagner-Meerwein rearrangement, where a hydrogenalkyl 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 theCope rearrangement.[53]
Cyclic rearrangements include cycloadditions and, more generally, pericyclic reactions, 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.[54]
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 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.[55] Because of the orbital character, the potential for developing stereoisomeric products upon cycloaddition is limited, as described by the Woodward–Hoffmann rules.[56]

## Biochemical reactions

Illustration of the induced fit model of enzyme activity
Biochemical reactions are mainly controlled by enzymes. These proteins can specifically catalyze a single reaction, so that reactions can be controlled very precisely. 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.[57]
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 proteins and carbohydrates from smaller units.[58]Bioenergetics studies the sources of energy for such reactions. An important energy source is glucose, which can be produced by plants via photosynthesis or assimilated from food. All organisms use this energy to produce adenosine triphosphate (ATP), which can then be used to energize other reactions.

## Applications

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 and mineral ores. It is essential to make the reaction as efficient as possible, maximizing the yield and minimizing the amount of reagents, energy inputs and waste. Catalysts are especially helpful for reducing the energy required for the reaction and increasing its reaction rate.[59][60]
Some specific reactions have their niche applications. For example, the thermite reaction is used to generate light and heat inpyrotechnics and welding. Although it is less controllable than the more conventional oxy-fuel weldingarc welding and flash welding, it requires much less equipment and is still used to mend rails, especially in remote areas.[61]

## 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 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 redistribution of substances in the human body. Faster reactions are usually studied with ultrafast laser spectroscopy where utilization of femtosecond lasers allows short-lived transition states to be monitored at time scaled down to a few femtoseconds.[62]