Reaction intermediates Knowledge-Based Study

Study notes

1. What is a Reaction Intermediate?

   Most chemical transformations do not take place in a single collision. Instead, they occur through a sequence of discrete, elementary steps known collectively as a reaction mechanism.

   A reaction intermediate is a chemical species that is produced in one elementary step and subsequently consumed in a later elementary step of the reaction mechanism.

   Key characteristics of reaction intermediates include:

   • They do not appear in the overall balanced chemical equation because their production and consumption cancel out.
   • They have real, finite lifetimes and exist as distinct chemical entities with fully formed, stable chemical bonds (unlike transition states).
   • Under suitable conditions, some intermediates can be observed, trapped, or characterized spectroscopically.
   • They typically have higher potential energy and lower stability than either the initial reactants or the final products, making them relatively transient.

   Potential Energy Profiles and Intermediates

   A reaction coordinate diagram visually represents the energy changes that occur along the reaction pathway.

   When examining a multi-step reaction profile:

   • Transition States (Activated Complexes): Correspond to energy maxima (peaks) along the pathway. They represent fleeting arrangements of atoms with partially formed or broken bonds at the highest energy configuration of each elementary step. A transition state has no measurable lifetime (on the order of femtoseconds) and cannot be isolated.
   • Reaction Intermediates: Correspond to local energy minima (troughs or shallow valleys) positioned between two consecutive transition states. Because an intermediate resides in a potential energy well, it requires activation energy to proceed forward or to revert backward.

   In a mechanism with $n$ elementary steps, there are $n$ transition states and $n - 1$ reaction intermediates.

   Intermediate versus Catalyst

   Both reaction intermediates and catalysts do not appear in the overall balanced chemical equation, which can lead to confusion. However, their roles and appearances in the mechanism are fundamentally opposite:

   • Reaction Intermediate: Formed first (as a product of an early step) and consumed later (as a reactant in a subsequent step).
     Sequence: Reactants $ightarrow$ Intermediate $ightarrow$ Products
   • Catalyst: Added initially (as a reactant in an early step) and regenerated later (as a product of a subsequent step).
     Sequence: Catalyst consumed $ightarrow$ Catalyst regenerated

   Illustrative Mechanism:

   Consider the decomposition of ozone catalyzed by chlorine atoms:

   Step 1: $\text{Cl}(g) + \text{O}_3(g) ightarrow \text{ClO}(g) + \text{O}_2(g)$

   Step 2: $\text{ClO}(g) + \text{O}(g) ightarrow \text{Cl}(g) + \text{O}_2(g)$

   Overall Reaction: $\text{O}_3(g) + \text{O}(g) ightarrow 2\,\text{O}_2(g)$

   In this mechanism:

   • $\text{Cl}(g)$ is a catalyst because it is introduced as a reactant in Step 1 and regenerated as a product in Step 2.
   • $\text{ClO}(g)$ is a reaction intermediate because it is generated in Step 1 and consumed in Step 2.

   Handling Intermediates in Experimental Rate Laws

   An experimentally determined rate law can only be expressed in terms of measurable quantities: reactants, overall products, or catalysts. Intermediates cannot be part of the final rate law because their concentrations are typically tiny, transient, and difficult to measure directly.

   The Steady-State Approximation (SSA)

   When an intermediate is highly reactive, its concentration remains low and virtually constant throughout the main course of the reaction. This principle is called the steady-state approximation:

   $$\frac{d[\text{Intermediate}]}{dt} = \text{Rate of formation} - \text{Rate of consumption} \approx 0$$

   By setting the net rate of change of the intermediate to zero, you can algebraically express the intermediate's concentration in terms of known reactant concentrations and rate constants, allowing substitution into the overall rate law expression.