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18.12: Reaction Intermediate

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  • Ozone \(\left( \ce{O_3} \right)\) depletion in the atmosphere is of significant concern. This gas serves as a protection against the ultraviolet rays of the sun. Ozone is naturally depleted in addition to the depletion caused by human-made chemicals. The depletion reaction is a two-step process:

    \[\ce{O_3} + \text{ultraviolet light} \rightarrow \ce{O_2} + \ce{O} \cdot \: \text{(free radical) slow reaction}\]

    \[\ce{O} \cdot + \ce{O_3} \rightarrow 2 \ce{O_2} \: \text{fast reaction}\]

    The free radical is not a part of the overall equation, but can be detected in the lab.


    Reaction mechanisms describe how the material in a chemical reaction gets from the initial reactants to the final products. One reaction that illustrates a reaction mechanism is the reaction between nitrogen monoxide and oxygen to form nitrogen dioxide:

    \[2 \ce{NO} \left( g \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{NO_2} \left( g \right)\]

    It may seem as though this reaction would occur as the result of a collision between two \(\ce{NO}\) molecules with one \(\ce{O_2}\) molecule. However, careful analysis of the reaction has detected the presence of \(\ce{N_2O_2}\) during the reaction. A proposed mechanism for the reaction consists of two elementary steps:

    Step 1: \(2 \ce{NO} \left( g \right) \rightarrow \ce{N_2O_2} \left( g \right)\)

    Step 2: \(\ce{N_2O_2} \left( g \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{NO_2} \left( g \right)\)

    In the first step, two molecules of \(\ce{NO}\) collide to form a molecule of \(\ce{N_2O_2}\). In the second step, that molecule of \(\ce{N_2O_2}\) collides with a molecule of \(\ce{O_2}\) to produce two molecules of \(\ce{NO_2}\). The overall chemical reaction is the sum of the two elementary steps:

    \[\begin{align} 2 \ce{NO} \left( g \right) &\rightarrow \cancel{\ce{N_2O_2} \left( g \right)} \\ \cancel{\ce{N_2O_2} \left( g \right)} + \ce{O_2} \left( g \right) &\rightarrow 2 \ce{NO_2} \left( g \right) \\ \hline 2 \ce{NO} \left( g \right) + \ce{O_2} \left( g \right) &\rightarrow 2 \ce{NO_2} \left( g \right) \end{align}\]

    The \(\ce{N_2O_2}\) molecule is not part of the overall reaction. It was produced in the first elementary step, then reacts in the second elementary step. An intermediate is a species which appears in the mechanism of a reaction, but not in the overall balanced equation. An intermediate is always formed in an early step in the mechanism and consumed in a later step.

    Figure 18.12.1: Nitrogen dioxide (left) and dinitrogen tetroxide (right).


    • The role of intermediates in reaction mechanisms is described.


    • CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.