Problem: The proposed mechanism for a given reaction is as follows:Step One: 2 NO (g)    N2O2 (g)Step Two: H2(g) + N2O2(g)    H2O(g) + N2O(g)      Rate - determining StepStep Three: N2O(g) + H2(g)    N2(g) + H2O(g)    FastUnder what conditions would this reaction be first-order in H2? Please answer and justify your reasoning in 1-2 complete sentences on the line provided.

FREE Expert Solution

For this problem, we are being asked on what conditions would this reaction be first order in H2.

In this case, we have to revisit the steady-state approximation for the intermediates 

Recall that the rate law only focuses on the reactant concentrations and has a general form of:

rate law=kAxBy

k = rate constant
A & B = reactants
x & y = reactant orders

  • A steady-state approach makes use of the assumption that the rate of production of an intermediate is equal to the rate of its consumption.
  • We will be starting this with:

Rate law = -d[H2]dt= k2[H2][N2O2] (1)

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Problem Details

The proposed mechanism for a given reaction is as follows:

Step One: 2 NO (g)    N2O2 (g)

Step Two: H2(g) + N2O2(g)    H2O(g) + N2O(g)      Rate - determining Step

Step Three: N2O(g) + H2(g)    N2(g) + H2O(g)    Fast

Under what conditions would this reaction be first-order in H2? Please answer and justify your reasoning in 1-2 complete sentences on the line provided.

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