Chemical Kinetics

 60) 2NO(g) + H₂(g) N₂O(g) + H₂O(g)

       NO(g) + NO(g) N₂O₂(g)

       N₂O₂(g) + H₂(g) N₂O(g) + H₂O(g)

       (a) NO(g) + NO(g) N₂O₂(g)                      (Step 1)

             N₂O₂(g) + H₂(g) N₂O(g) + H₂O(g)     (Step 2)
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             2NO(g) + H₂(g) N₂O(g) + H₂O(g)

       (b) Rate1 = k[NO][NO] = k[NO]²

             Rate2 = k[N₂O₂][H₂]

       (c) N₂O₂(g) is the intermediate because it is formed in one elementary step and consumed in the next. An intermediate is neither a reactant nor a product in the overall reaction.

       (d) Rate = k[NO]²[H₂]

             Because [H₂] exists in the rate law, the second step must be the slow (rate-determining step). The first step must be the fast step.