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A reaction intermediate or an intermediate is a molecular entity that is formed from the reactants (or preceding intermediates) and reacts further to give the directly observed products of a chemical reaction. Most chemical reactions are stepwise, that is they take more than one elementary step to complete. An intermediate is the reaction product of each of these steps, except for the last one, which forms the final product. Reactive intermediates are usually short lived and are very seldom isolated. Also, owing to the short lifetime, they do not remain in the product mixture.
For example, consider this hypothetical stepwise reaction:
- A + B → C + D
The reaction includes these elementary steps:
- A + B → X*
- X* → C + D
The chemical species X* is an intermediate.
The IUPAC Gold Book defines an intermediate as a molecular entity (atom, ion, molecule...) with a lifetime appreciably longer 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, and thus, intermediates correspond to potential energy minima of depth greater than available thermal energy arising from temperature, (RT, where R is gas constant and T is temperature).
Many intermediates are short-lived and highly reactive, thus having a low concentration in the reaction mixture. As is always the case when discussing chemical kinetics, definitions like fast/slow short/long-lived are relative, and depend on the relative rates of all the reactions involved. 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.
Methane chlorination is a chain reaction. If only analyze products and reactants are analyzed, the result is:
However, this reaction has intermediates. A sequence of irreversible first-order reactions is taken place until we arrive at the final product.
CH4→ CH3Cl→ CH2Cl2→ CHCl3→ CCl4
Reactants: CH4 Products: CCL
The other elements are reaction intermediates: CH3Cl, CH2Cl2 and CHCl3.
These are the set of irreversible first-order reactions:
CH4 - CH3Cl
CH3Cl - CH2Cl2
CH2Cl2 - CHCl3
CHCl3 - CCl4
The elements’ concentration can be calculated by integrating the system of kinetic equations.
Example: hydrocarbon chlorination (methane) If only products and reactants are analyzed the reaction is: CH4 → CCl4.
However, this reaction has intermediate reactants which are formed during a sequence of irreversible reactions until we arrive at the final product. This reaction takes place in 4 steps. This is why it’s called a chain reaction.
Initiation: This reaction can occur by thermolysis (heating) or photolysis (absorption of light) leading to the breakage of a molecular chlorine bond.
Cl-Cl → Cl. + Cl.
When the bond is broken it ends up to two highly reactive chlorine molecules.
Propagation: It has two parts. As chlorine alone is unstable, it reacts with methane’s hydrogen. At the end, one molecule of hydrochloric acid and one molecule of a radical methane are gotten. The bond of a second Cl2 molecule is broken. And one of these radicals binds to the protonated methane. At this moment there are several compounds which will be stabilized at termination step.
Termination: It takes place when the reactants have run out. This step stabilizes all the reactants created during the reaction. The radicals combine and the chain carriers are lost.
There are different combinations: Union of methane radicals from a C-C bond leading to ethane.
CH3. + CH3. → CH3 - CH3
Union of one methane radical to a Cl radical forming chloromethane.
CH3. + Cl. → CH3 – Cl
Union of two Cl radicals to form chlorine gas.
Cl. + Cl. → Cl-Cl
- 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]