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How can the reaction between hydrogen and iodine chloride occur in two steps?

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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.

In chemical kinetics, the reaction between hydrogen and iodine chloride ( $H_{2} + 2 I C l \to I_{2} + 2 H C l$ ) is considered a composite reaction because its experimental rate law ( $v = k [H_{2}] [I C l]$ ) does not match the stoichiometry of the balanced equation.1

This reaction occurs in two consecutive steps involving an intermediate:1

Step 1 (Slow): A single hydrogen molecule reacts with one molecule of iodine chloride to form hydrogen iodide and hydrogen chloride:
$H_{2} + I C l \to H I + H C l$
Step 2 (Fast): The hydrogen iodide formed in the first step reacts rapidly with a second molecule of iodine chloride to produce the final products:
$H I + I C l \to H C l + I_{2}$

Because the second step is fast, the hydrogen iodide is consumed as quickly as it is produced. Consequently, the overall reaction rate is determined solely by the rate of the first, slower step, which explains why the reaction is second-order overall rather than third-order.1