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Chemistry LibreTexts

18.14: Rate-Determining Step

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    Rate-Determining Step

    The determination of a reaction mechanism can only be made in the laboratory. When a reaction occurs in a sequence of elementary steps, the overall reaction rate is governed by whichever one of those steps is the slowest. The rate-determining step is the slowest step in the sequence of steps in a reaction mechanism. To get an idea of how one step is rate determining, imagine driving on a one-lane road where it is not possible to pass another vehicle. The rate of flow of traffic on such a road would be dictated by whatever car is traveling at the lowest speed. The decomposition of hydrogen peroxide is discussed below and illustrates how reaction mechanisms can be determined through experimental studies.

    Decomposition of Hydrogen Peroxide

    Recall that a catalyst is a substance which increases the rate of a chemical reaction without being consumed. Catalysts lower the overall activation energy for a reaction by providing an alternative mechanism for the reaction to follow. One such catalyst for the decomposition of hydrogen peroxide is iodide ions \(\left( \ce{I^-} \right)\).

    \[2 \ce{H_2O_2} \left( aq \right) \overset{\ce{I^-}}{\rightarrow} 2 \ce{H_2O} \left( l \right) + \ce{O_2} \left( g \right)\]

    By experiment, the rate of reaction is found to be first-order with respect to both \(\ce{H_2O_2}\) and \(ce{I^-}\) and second order overall.

    \[\text{rate} = k \left[ \ce{H_2O_2} \right] \left[ \ce{I^-} \right]\]

    The reaction cannot occur in one step corresponding to the overall balanced equation. If it did, the reaction would be second-order with respect to \(\ce{H_2O_2}\) since the coefficient of the \(\ce{H_2O_2}\) in the balanced equation is a 2. A reaction mechanism can be constructed which accounts for the rate law and for the detection of the \(\ce{IO^-}\) ion as an intermediate. It consists of two bimolecular elementary steps.

    Step 1: \(\ce{H_2O_2} \left( aq \right) + \ce{I^-} \left( aq \right) \rightarrow \ce{H_2O} \left( l \right) + \ce{IO^-} \left( aq \right)\)

    Step 2: \(\ce{H_2O_2} \left( aq \right) + \ce{IO^-} \left( aq \right) \rightarrow \ce{H_2O_2} \left( l \right) + \ce{O_2} \left( g \right) + \ce{I^-} \left( aq \right)\)

    If step 2 is the rate-determining step, then the rate law for that step will be the rate law for the overall reaction.

    \[\text{rate} = k \left[ \ce{H_2O_2} \right] \left[ \ce{I^-} \right]\]

    The rate law for the slow step of the proposed mechanism agrees with the overall experimentally determined rate law. The \(\ce{IO^-}\) is present as an intermediate in the reaction. The iodide catalyst also appears in the mechanism. It is consumed in the first elementary step and then is regenerated in the second step. That is the requirement for a catalyst - that is that it is not used up in the reaction.

    Summary

    • The rate-determining step in a reaction is defined.
    • The process for determining the rate-determining step is described.

    Contributors

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