Reaction intermediate

In chemistry, a reaction intermediate or an intermediate is a molecular entity that is formed from the reactants (or preceding intermediates) but is consumed in further reactions in stepwise chemical reactions that contain multiple elementary steps. Intermediates are the reaction product of one elementary step, but do not appear in the chemical equation for an overall chemical equation. ^[1]

For example, consider this hypothetical stepwise reaction:

{\ce {A + B -> C + D}}

The reaction includes two elementary steps:

{\ce {A + B -> X}}

{\ce {X -> C + D}}

In this example, X is a reaction intermediate.

IUPAC Definition

The IUPAC Gold Book^[2] defines an intermediate as a compound that has a lifetime greater than a molecular vibration that is formed (directly or indirectly) from the reactants and reacts further to give (either directly or indirectly) the products of a chemical reaction. The lifetime condition distinguishes true, chemically distinct intermediates from vibrational states or such transition states which, by definition have lifetimes close to that of molecular vibration.

Kinetically, intermediates are often consumed quickly in a step-wise mechanism. The designation of "fast" or "slow" consumption speed is relative, and a relative intermediate are sometimes separated from a reaction intermediate based on being relatively short-lived. Reactive intermediates are an unstable type of reaction intermediate, and are usually short-lived, high-energy, and seldom isolated. They do not remain in the product mixture due to their short lifetime, in contrast to other reaction intermediates.

Species that are short-lived in one reaction mechanism, can be considered stable in others and molecular entities that are intermediates in some mechanisms can be stable enough to be detected, identified, isolated or used as reactants in (or be the products of) other reactions. Reaction intermediates are often free radicals or unstable ions. Oxidizing radicals (OOH and OH) found in combustion reactions are so reactive that a high temperature is required to constantly produce them, in order to compensate their disappearance, or the combustion reaction will cease.

When the necessary conditions of the reaction no longer prevail, these intermediates react further and no longer remain in the reaction mixture. There are some operations where multiple reactions are run in the same batch. For example, in an esterification of a diol, a monoester product is formed first, and may be isolated, but the same reactants and conditions promote a second reaction of the monoester to a diester. The lifetime of such an "intermediate" is considerably longer than the lifetime of the intermediates of the esterification reaction itself (the tetrahedral intermediate).

Chemical processing industry

In the chemical industry, the term intermediate may also refer to the (stable) product of a reaction that is itself valuable only as a precursor chemical for other industries. A common example is cumene which is made from benzene and propylene and used to make acetone and phenol in the cumene process. The cumene itself is of relatively little value in and of itself, and is typically only bought and sold by chemical companies.

Example

Methane chlorination

Methane chlorination is a chain reaction. If only the products and reactants are analyzed, the result is:

${\ce {CH4 + 4Cl2->CCl4 +4HCl}}$

However, this reaction has 3 intermediate reactants which are formed during a sequence of 4 irreversible second order reactions until we arrive at the final product. This is why it's called a chain reaction. Following only the carbon containing species in series:

${\ce {CH4 -> CH3Cl -> CH2Cl2 -> CHCl3 -> CCl4}}$

Reactants: ${\ce {CH4 + 4Cl2}}$

Products: ${\ce {CCl4 + 4HCl}}$

The other species are reaction intermediates: ${\ce {CH3Cl, CH2Cl2, CHCl3}}$

These are the set of irreversible second-order reactions:

${\ce {CH4 + Cl2->CH3Cl + HCl}}$

${\ce {CH3Cl + Cl2->CH2Cl2 + HCl}}$

${\ce {CH2Cl2 + Cl2->CHCl3 + HCl}}$

${\ce {CHCl3 + Cl2->CCl4 + HCl}}$

These intermediate species’ concentrations can be calculated by integrating the system of kinetic equations. The full reaction is a free radical propagation reaction which is filled out in detail below.

Initiation: This reaction can occur by thermolysis (heating) or photolysis (absorption of light) leading to the breakage of a molecular chlorine bond. ${\ce {Cl-Cl ->[h \nu] Cl. + Cl.}}$

When the bond is broken it produces two highly reactive chlorine atoms.

Propagation: This stage has two distinct reaction classes. The first is the stripping of a hydrogen from the carbon species by the chlorine radicals. This occurs because chlorine atoms alone are unstable, and these chlorine atoms react with one the carbon species' hydrogens. The result is the formation of hydrochloric acid and a new radical methyl group.

${\ce {CH3-H + Cl. -> CH3. + H-Cl}}$

${\ce {CH2Cl-H + Cl. -> CH2Cl. + H-Cl}}$

${\ce {CHCl2-H + Cl. -> CHCl2. + H-Cl}}$

${\ce {CCl3-H + Cl. -> CCl3. + H-Cl}}$

These new radical carbon containing species now react with a second Cl₂ molecule. This regenerates the chlorine radical and the cycle continues. This reaction occurs because while the radical methyl species are more stable than the radical chlorines, the overall stability of the newly formed chloromethane species more than makes up the energy difference.

${\ce {CH3. + Cl-Cl -> CH3Cl + Cl.}}$

${\ce {CH2Cl. + Cl-Cl -> CH2Cl2 + Cl.}}$

${\ce {CHCl2. + Cl-Cl -> CHCl3 + Cl.}}$

${\ce {CCl3. + Cl-Cl -> CCl4 + Cl.}}$

During the propagation of the reaction, there are several highly reactive species that will be removed and stabilized at the termination step.

Termination: This kind of reaction takes place when the radical species interact directly. The products of the termination reactions are typically very low yield in comparison to the main products or intermediates as the highly reactive radical species are in relatively low concentration in relation to the rest of the mixture. This kind of reaction produces stable side products, reactants, or intermediates and slows the propagation reaction by lowering the number of radicals available to propagate the chain reaction.

There are many different termination combinations, some examples are:

Union of methyl radicals from a C-C bond leading to ethane (a side product).

${\ce {CH3. + CH3. -> CH3-CH3}}$

Union of one methyl radical to a Cl radical forming chloromethane (another reaction forming an intermediate).

${\ce {CH3. + Cl. -> CH3Cl}}$

Union of two Cl radicals to reform chlorine gas (a reaction reforming a reactant).

${\ce {Cl. + Cl. -> Cl2}}$

References

^ Moore, John W. (2015). Chemistry : the molecular science. Conrad L. Stanitski (Fifth edition ed.). Stamford, CT. ISBN 978-1-285-19904-7. OCLC 891494431. {{cite book}}: |edition= has extra text (help)CS1 maint: location missing publisher (link)
^ IUPAC Goldbook definition of intermediate

Francis A. Carey; Richard J. Sundberg (1985). Advanced organic chemistry Structure and mechanisms. ISBN 978-0-306-41198-4.^{[page needed]}
March, Jerry (1985). Advanced Organic Chemistry Reactions, Mechanisms, and Structure. John Wiley & Sons. ISBN 978-0-471-85472-2.^{[page needed]}
"Cloración del metano".
"Write notes on consecutive reactions with example class 10 chemistry CBSE".

[1] Moore, John W. (2015). Chemistry : the molecular science. Conrad L. Stanitski (Fifth edition ed.). Stamford, CT. ISBN 978-1-285-19904-7. OCLC 891494431. {{cite book}}: |edition= has extra text (help)CS1 maint: location missing publisher (link)

[2] IUPAC Goldbook definition of intermediate

[1]

[2]

IUPAC Definition

Chemical processing industry

Example

See also

References