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Why can some chemical reactions occur on femtosecond timescales?

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A graph showing a potential energy curve for a chemical reaction, with the y-axis labeled potential energy and the x-axis labeled progress of reaction. The curve rises from the initial state (reactants) to a peak labeled transition state (activated complex), indicating activation energy, then drops to the final state (products). — chemical kinetics In the Arrhenius equation, the activation energy (E) represents the minimum amount of energy required to transform reactants into products in a chemical reaction. On a potential energy curve, the value of the activation energy is equivalent to the difference in potential energy between particles in an intermediate configuration (known as the activated complex, or transition state) and particles of reactants in their initial state. The activation energy thus can be visualized as a barrier that must be overcome by reactants before products can be formed.

Chemical reactions occur on femtosecond timescales because they involve fundamental physical processes, such as the redistribution of energy and the breaking or formation of chemical bonds, which take place at these extremely rapid intervals.1

While a nanosecond is sufficient to observe changes in chemical identity, the femtosecond range is required to capture the underlying physical dynamics of a reaction. For instance, the time required for a chemical bond length to change by $10^{- 10}$ meters can be as short as approximately 100 femtoseconds. These processes are studied using pulse and probe methods, such as flash-photolysis, where high-intensity, short-duration light pulses allow scientists to observe these transient atomic and molecular movements.1

In the context of chemical kinetics, these observations provide insight into the reaction mechanism and the atomic motions occurring along the reaction path, helping to validate theories like transition-state theory and molecular dynamics.213