Reaction Mechanisms: Understanding the Steps Behind Chemical Reactions
What Are Elementary Steps?
Elementary step
An elementary step is a single molecular event in a reaction mechanism where reactants are converted into products or intermediates.
Each step involves a specific collision or transformation of particles, and it cannot be broken down further.
- Consider the reaction:
$$
2NO_2(g) + F_2(g) \to 2NO_2F(g)
$$ - This reaction may occur in two elementary steps: $$NO_2(g) + F_2(g) \to NO_2F(g) + F(g)$$ $$NO_2(g) + F(g) \to NO_2F(g)$$
- Here, the overall reaction is the sum of these steps, but each step represents a distinct molecular interaction.
- $F(g)$ is an intermediate: it is produced in Step 1 and consumed in Step 2.
- Intermediates are crucial for understanding the step-by-step nature of the reaction.
Molecularity of Elementary Steps
Molecularity
The molecularity of an elementary step refers to the number of reacting particles (atoms, ions, or molecules) that must collide simultaneously to drive a chemical change.
- Unimolecular: A single molecule decomposes or rearranges (e.g., $A \to B + C$).
- Bimolecular: Two particles collide and react (e.g., $A + B \to C$).
- Termolecular: Three particles collide simultaneously (e.g., $A + B + C \to D$). These are rare due to the low probability of three particles colliding at the same time.
Unlike molecularity, which applies to elementary steps, reaction order is determined experimentally and applies to the overall reaction.
The Rate-Determining Step: A Bottleneck for Reaction Rates
In a multistep reaction, not all steps occur at the same speed.
Rate-determining step
The slowest step in the mechanism is called the rate-determining step (RDS).
This step acts as a bottleneck, limiting the overall reaction rate, much like how the slowest person in a relay race determines the team's overall time.
Why Does the RDS Matter?
- The RDS has the highest activation energy ($E_a$) among all the steps.
- Since the rate of a reaction depends on the energy barrier that must be overcome, the RDS dictates the overall reaction rate.
- If Step 1 is slow and Step 2 is fast, the overall reaction rate depends only on Step 1.
- If Step 2 is slow, even if Step 1 is fast, the reaction cannot proceed faster than Step 2.
Identifying Intermediates and Transition States
What Are Intermediates?
Intermediate
An intermediate is a species that is formed in one elementary step and consumed in another. It is not present in the overall reaction equation because it does not appear in the final products or reactants.
- In the following reaction mechanism: $$NO_2(g) + F_2(g) \to NO_2F(g) + F(g)$$ $$NO_2(g) + F(g) \to NO_2F(g)$$
- The $F(g)$ species is an intermediate: it is produced in Step 1 and consumed in Step 2.
Intermediates are often unstable and exist only briefly during the reaction.
What Are Transition States?
Transition state
A transition state represents the highest-energy arrangement of atoms during an elementary step. It is the point where bonds are partially broken and formed, and the system is at the peak of its energy profile.
Unlike intermediates, transition states cannot be isolated: they exist only momentarily as the reaction progresses.
Visualizing energy profiles can help you distinguish between intermediates (valleys) and transition states (peaks).
Evaluating Proposed Mechanisms
- Chemists propose reaction mechanisms based on experimental evidence.
- To determine whether a mechanism is plausible, it must meet the following criteria:
1. Consistency with the Overall Reaction
The sum of all elementary steps in the mechanism must match the stoichiometry of the overall reaction.
- Overall reaction: $$2NO_2(g) + F_2(g) \to 2NO_2F(g)$$
- Mechanism: $$NO_2(g) + F_2(g) \to NO_2F(g) + F(g)$$ $$NO_2(g) + F(g) \to NO_2F(g)$$
- Adding these steps gives the overall reaction.
2. Agreement with Experimental Rate Data
The rate law derived from the mechanism must match the experimentally determined rate law.
- Experimental rate law: $\nu = k[NO_2][F_2]$
- Mechanism:
- Step 1 (slow): $$NO_2(g) + F_2(g) \to NO_2F(g) + F(g)$$
- Step 2 (fast): $$NO_2(g) + F(g) \to NO_2F(g)$$
- Since Step 1 is the RDS, the rate law for the overall reaction is determined by this step: $\nu = k[NO_2][F_2]$, which matches the experimental data.
- Do not assume the rate law can be deduced from the overall reaction equation.
- It must be determined experimentally or derived from the RDS.
Evaluating a Mechanism
- Consider the reaction:
$$
2NO(g) + Cl_2(g) \to 2NOCl(g)
$$ - Experimental rate law: $\nu = k[NO]^2[Cl_2]$
- Proposed mechanism:
- $NO(g) + Cl_2(g) \to NOCl_2(g)$ (slow)
- $NOCl_2(g) + NO(g) \to 2NOCl(g)$ (fast)
Step-by-Step Evaluation:
Overall Reaction:
- Adding the steps gives:
$$
NO(g) + Cl_2(g) + NOCl_2(g) + NO(g) \to NOCl_2(g) + 2NOCl(g)
$$ - Canceling $NOCl_2(g)$ (intermediate), we get:
$$
2NO(g) + Cl_2(g) \to 2NOCl(g)
$$
This matches the overall reaction.
Rate Law:
- Since Step 1 is the RDS, the rate law is determined by this step:
$$
\nu = k[NO][Cl_2]
$$ - However, the experimental rate law is $\nu = k[NO]^2[Cl_2]$.
- This suggests that the proposed mechanism is incorrect.
- Why are transition states considered the "peaks" of energy profiles, while intermediates are "valleys"?
- How might experimental techniques, such as spectroscopy, help identify intermediates in a reaction mechanism?
- Why are termolecular reactions rare, and how does this affect the design of industrial chemical processes?
