Intermediates

Definition

Intermediates are species that appear in the steps of a reaction mechanism but not in the overall balanced equation. They are formed in one elementary step and consumed in a subsequent step.

5 Must Know Facts For Your Next Test

1. Intermediates are typically unstable and have a short lifetime.
2. They do not appear in the overall reaction equation because they cancel out when combining elementary steps.
3. The presence of intermediates can be inferred from experimental data, such as reaction rate studies.
4. Intermediates help explain the mechanism by which reactants are converted to products.
5. In energy diagrams, intermediates correspond to valleys between transition states.

Review Questions

• What is an intermediate in a chemical reaction?
• Why don't intermediates appear in the overall balanced equation for a reaction?
• How can you identify an intermediate in a reaction mechanism?

Related terms

Elementary Step: A single reaction event with its own specific rate law and molecularity within a larger mechanism.

Transition State:

A high-energy state during a chemical reaction where bonds are partially broken and formed, preceding the formation of an intermediate or product.

Rate-Determining Step:

The slowest step in a reaction mechanism that determines the overall reaction rate.