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An intermediate is a short-lived species formed in the middle of a reaction pathway. It is formed from the reactants and has a discrete, independent existence before it reacts further to yield the final products. Intermediates are often unstable and have a short lifetime, but they can be identified and characterized in some cases.

An activated complex, also known as the transition state, is the highest-energy state along the reaction pathway. This is the point at which the old bonds are breaking, and the new bonds are forming. The activated complex is not an isolated species; rather, it represents an unstable arrangement of atoms during the transition from reactants to products. It cannot be isolated or directly observed experimentally but can be studied indirectly using computational methods and transition state theory.

1. Stability: An intermediate is a relatively more stable species compared to an activated complex. The activated complex represents a high-energy transition state that lies at the top of the energy profile. 2. Existence: An intermediate has a discrete, independent existence, although short-lived, while an activated complex is not an isolated species and cannot be isolated or directly observed. 3. Role in reaction: Intermediates are formed and consumed during a reaction pathway, while an activated complex represents the transition state at which old bonds break and new bonds form. 4. Characterization: Intermediates can sometimes be identified and characterized, whereas the activated complex is studied indirectly through computational methods and transition state theory. In conclusion, an intermediate is a short-lived species formed during a reaction, while an activated complex represents the transition state between the reactants and products. The main differences between the two include stability, existence, their role in the reaction, and methods of characterization.