Final answer:
A proposed reaction mechanism for the decomposition of N₂O₅ into NO₂ and O₂ involves multiple steps but does not match the provided information. Upon revising the steps provided, the reaction proceeds with NO forming an intermediate N₂O₂, which then reacts with O₂ to give NO₂.
Step-by-step explanation:
The reaction 2N₂O₅(g) ⟷ 4NO₂(g) + O₂(g) involves the decomposition of dinitrogen pentoxide (N₂O₅) into nitrogen dioxide (NO₂) and dioxygen (O₂), which can be explained through a proposed mechanism involving multiple steps. Breaking down complex reactions into elementary steps is fundamental in understanding the rate law and mechanism of a reaction.
The mechanism for this reaction can be suggested in multiple steps. Step 1: involves the formation of NO₂ and a nitrate radical NO₃, and can be represented as N₂O₅(g) ⇌ NO₂(g) + NO₃(g). Step 2: NO₃ might further react with NO₂ to generate N₂O₄, and subsequently, N₂O₄ breaks down to form more NO₂. Since these steps do not match the information provided, let's consider the steps you've provided.
The first step 2NO(g) → N₂O₂(g), involves the dimerization of two NO molecules into N₂O₂. The second step is the reaction of this dimer, N₂O₂, with O₂ to produce two molecules of NO₂: N₂O₂(g) + O₂(g) → 2NO₂(g). These steps form an overall reaction that matches the original equation by summing the individual reactions. The reaction rate expressions depend on the reaction's stoichiometry and its mechanism.
Using the provided steps, the mechanism indicates a chain of reactions where NO and O₂ are involved in the formation of the intermediate N₂O₂, which then reacts to produce the NO₂ in the final step. The rate-determining step, which controls the overall rate of the reaction, appears to be the second step involving the conversion of N₂O₂ to NO₂ due to the presence of O₂. This mechanism helps to explain the observed kinetics and behavior of the reaction under study.