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Why does the steady-state approximation not apply during an explosion?

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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.

The steady-state approximation is a method used in chemical kinetics that assumes the concentration of reaction intermediates remains small and effectively constant (zero rate of change) throughout a reaction. This approximation is valid only when intermediates are present at low, stable concentrations.1

The steady-state approximation does not apply during an explosion because the concentration of intermediates does not remain steady. Instead, explosions often involve a branching chain mechanism, where a single chain carrier reacts to produce multiple new chain carriers. This causes the concentration of intermediates to rise rapidly to high values, leading to an uncontrolled, accelerating reaction. Because the fundamental requirement of the steady-state treatment—that the intermediate concentration remains small and constant—is violated, the approximation cannot be used to describe the kinetics of an explosion.1