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.