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## Reaction mechanism

Chemical reactions often take place in multiple steps and it’s not always obvious from looking at the chemical equation. For example, the decomposition of ozone: From the equation, it appears that O₃ decomposes to O₂ in one step. In reality, it’s more likely to take place in the following sequence: The above shows the decomposition taking place with two steps, we call each step an elementary step. An elementary step is an individual step in the reaction mechanism that cannot be further broken down into smaller steps. Reactions that involve more than one elementary step are called complex reactions.

## Intermediates

In the example above, the oxygen atom “O” produced in the first step of the mechanism is consumed in the second step. The entity is known as a “reaction intermediate” and it does not appear in the final chemical reaction.

## Rate determining step

How fast can a reaction reach equilibrium or completion? The rate of the reaction is dictated by the slowest elementary step of the reaction. This step is called the “rate-determining step“. No matter how fast the speed of the other elementary steps are, the reaction cannot go faster than the slowest step. The diagram above is called a reaction coordinate diagram or energy profile diagram. It shows how the energy of the system changes during a chemical reaction.

The rate-determining step is the step with the largest activation energy. In the diagram above, the 1st elementary step is the rate-determining step since it requires more energy for it to proceed forward.

## Molecularity

What happens at the molecular level? Molecularity of a reaction tells you how many number of molecules or ions are coming together in a single step reaction to form products.

A unimolecular reaction is an elementary reaction that a single molecule is rearranged to form the product.

A → B

eg. PCl₅ → PCl₃ + Cl₂

Rate law = k[A]… first order reaction

A bimolecular reaction involves the collision of two molecules in an elementary reaction to form the product.

A + B → products … two molecules from two different reactants

NO + O₃ → NO₂ + O₂

Rate law = k[A][B]

2A  → products … two molecules of the same reactant

2HI → H₂ + I₂

Rate law = k[A]²

A termolecular reaction involves the collision of three molecules in an elementary reaction to form the product.

A + A + A → products

A + B + C → products

A + B + B → products

Rate law = k[A]³ … 3rd order

Rate law = k[A][B][C] … 1st order in A,B,C

Rate law = k[A][B]² … 1st order in A and 2nd order in B

Note: A termolecular reaction involving collision of three molecules with right energy and orientation simultaneously is very rare. It is more likely that each elementary step is a bimolecular reaction.

## Reaction mechanism

Suppose you have a reaction given by the equation:

NO₂(g) + CO(g)  → CO₂(g) + NO₃(g)

In an experiment (carried out above 225°C), the rate law determined for the reaction was:

rate = k•[NO₂][CO]

In an experiment (carried out below 225°C), the rate law obtained for the reaction was:

rate = k•[NO₂]²

Above 225°C, the reaction is first order with respect to NO₂ and CO. Below 225°C the reaction is second order with respect to NO₂. Suggest a mechanism that is consistent with the above rate laws.

## Proposing a mechanism

Above 225°C, we expect the reaction to happen at a faster rate and a possible mechanism we can suggest is the simple one-step bimolecular collision:

NO₂(g) + CO(g)  → CO₂(g) + NO₃(g)

Above 225°C, the proposed mechanism is consistent with the experimental rate law:

rate = k•[NO₂][CO]

Below 225°C, the reaction would happen at a slower rate and we can suggest a two-step mechanism:

Mechanism 1

NO₂(g) + NO₂(g)  → NO₃(g) + NO(g) … (slow step)

NO₃(g) + CO(g)  → NO₂(g) + CO₂(g) … (fast step)

The slowest step is the rate determining step.

rate = k • [NO₂]²

Mechanism 2

NO₂(g) + NO₂(g)  → NO₃(g) + NO(g) … (fast step)

NO₃(g) + CO(g)  → NO₂(g) + CO₂(g) … (slow step)

The slowest step is the rate determining step.

rate = k • [NO₃][CO]

Conclusion

Below 225°C, using mechanism 1, the proposed mechanism is consistent with the experimental rate law:

rate = k • [NO₂]²

## Sum of individual steps

The sum of the individual steps must sum up to the overall equation. The intermediates do not appear as a final product. NO₂(g) + CO(g)  → CO₂(g) + NO₂(g)