Question

I'm really confused about those three concepts: radicals, reaction intermediate, activated complex.

1. First of all, in my course we have seen radicals in the context of chain radical reactions. But does it have a wider meaning ? (in older exams of my course I am seeing the word radical for reactions that are simple elementary reactions, no propagation and stuff).
2. I have seen the use of an asterisk in the notation of radicals in radical chain reactions. Is it a specific notation for the particular case of radicals inside chain reactions? Not only ?
3. Is it true that all radicals are reaction intermediate but not all reaction intermediate are radicals? The contrary ? None of the 2, how could we describe it then ?
4. About activated complex now, is there any way to distinguish an activated complex (which I understand represent a maximum of energy) from a classical reaction intermediate (which I understand represent a local minimum of energy) just by the look at the shape of the chemical reaction(s) ? (so without calculations using specific data about the components).
5. Can we calculate a pseudo-equilibrium constant (which is related to the fact that we have a steady state, correct me if I'm wrong) either in the case of complex activated and reaction intermediate ?
Thank you very much in advance !

EDIT: I understand your feeling that led you to close this topic. Tbh though, many students of our course were confused about those particular interactions and the point of the post was precisely to shed light on the links (or non links) between all those topics. I understand it may not be very meaningful to others who have had a different (and maybe clearer) learning structure (However, I suppose it would be possible that similar confusions could be happening elsewhere, not just with us). Also, I think it's worth mentioning that Poutnik's answers were very helpful and helped many of my classmates to whom I passed on the information. I respect your choice to close this post but at least you know how it happened behind the scenes.

Answer 1

Radicals are highly reactive species with unpaired electrons and can be reaction intermediates, but not all intermediates are radicals. An activated complex represents the highest energy state of reacting particles and can be distinguished from other intermediates on an energy profile. Pseudo-equilibrium constants are applicable to intermediates in a steady state but not to activated complexes.

Step-by-step explanation:

The concepts of radicals, reaction intermediates, and activated complexes are crucial in understanding the details of chemical reactions. Radicals are atoms, molecules, or ions with unpaired electrons that are very reactive. These may be involved in various types of reactions, including but not limited to radical chain reactions. The notation of a radical using an asterisk is not specific to chain reactions; it is a general notation representing the unpaired electron.

Indeed, all radicals can be considered as reaction intermediates because they appear transiently during the reaction process. However, not all reaction intermediates are radicals; intermediates can also be neutral or ionic species that are not radicals.

An activated complex, also known as a transition state, is an unstable arrangement of atoms at the peak of the energy barrier of the reaction. It represents a maximum in energy on the reaction coordinate and exists for a very short period (~10-13 seconds).

By looking at a reaction coordinate diagram, you can distinguish an activated complex from a classical reaction intermediate as the former corresponds to the highest energy point (energy maximum), while the latter are stable species that correspond to valleys (local minima) in the energy profile. Pseudo-equilibrium constants can be calculated for reaction intermediates in steady state situations but not typically for activated complexes due to their extremely short-lived nature.