2.11:
Energy Diagrams, Transition States, and Intermediates
Energy diagrams are used to represent the change in energy for the molecules involved in a chemical reaction.
The free energy is measured along the y-axis, and the reaction coordinate is plotted on the x-axis.
The reaction coordinate indicates the progress of the conversion of reactants to products. Peaks on the energy diagram represent the transition states, whereas the valleys represent the reactive intermediates.
As the reaction progresses, the reactants pass through an unstable state of maximum free energy, called the activated complex, or the transition state. They last for less than one picosecond and cannot be isolated. The reactant’s ability to achieve a given transition state depends on the value of the activation energy.
The double dagger symbol is used to describe transition states. If the transition state is easy to achieve, the associated delta-G double-dagger is small, and the reaction is fast. Contrarily, if the transition state is difficult to achieve, the associated delta-G double-dagger is large, and the reaction rate is slow.
According to Hammond's postulate, in a single-step exothermic process, the structure of the transition state closely resembles that of the reactants, as the transition state is closer in energy to the reactants than to the products.
Meanwhile, in a single-step endothermic process, the structure of the transition state closely resembles that of the products, as the transition state is closer in energy to the products than to the reactants.
A reactive intermediate corresponds to an energy minimum between two transition states. Highly reactive intermediates are difficult to isolate and have a short lifespan, while those with lower energies have longer lifetimes.
The most common reactive intermediates in organic chemistry involve carbon radicals or carbon centers without four bonds. For example, carbocations act as electron acceptors and carbanions act as electron donors.
2.11:
Energy Diagrams, Transition States, and Intermediates
Free-energy diagrams, or reaction coordinate diagrams, are graphs showing the energy changes that occur during a chemical reaction. The reaction coordinate represented on the horizontal axis shows how far the reaction has progressed structurally. Positions along the x-axis close to the reactants have structures resembling the reactants, while positions close to the products resemble the products. Peaks on the energy diagram represent stable structures with measurable lifetimes, while other points along the graph represent unstable structures that cannot be isolated.
This high-energy unstable structure is called the transition state or activated complex. In this high-energy process, bonds are in the process of being broken and/or formed simultaneously. The structure is so strained that it transitions into new, less strained structures.
George Hammond formulated a principle that relates the nature of a transition state to its location on the reaction diagram. The Hammond Postulate states that a transition state will be structurally and energetically similar to the species nearest to it on the reaction diagram. In the case of an exothermic reaction, the transition state resembles the reactant species, whereas, in the case of an endothermic reaction, the transition state resembles the products. In a multi-step reaction, each step has a transition state and corresponding activation energy. The transition states of such reactions are punctuated with reactive intermediates, which are represented as local minima on the energy diagrams.
Reactive intermediates are products of bonds breaking and cannot be isolated for prolonged periods of time. Some of the most common reactive intermediates in organic chemistry are carbon ions or radicals. Carbocations are electrophiles, and carbanions are nucleophiles. Carbon radicals have only seven valence electrons and may be considered electron deficient; however, they do not, in general, bond to nucleophilic electron pairs, so their chemistry exhibits unique differences from that of conventional electrophiles. Radical intermediates are often called free radicals.
Suggested Reading
- Brown, W.H., & Iverson, B.L., & Anslyn, V.E., & Foote S.C. (2014). Organic Chemistry. Mason, Ohio: Cengage Learning, 118-123.
- Klein, D. (2017). Organic Chemistry. New Jersey, NJ: Wiley, 183-188.