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)    FastUse the steady-state approximation on intermediates to write the rate law for this reaction in terms of  . Your law should include only elementary rate constants and species that appear in the overall reaction. Show all work to support your reasoning and answer. Note: you may Not assume step one is a fast pre-equilibrium.

FREE Expert Solution

For this problem, we have to write the rate law for this reaction in terms of -d[H₂]/dt using steady-state approximation

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

$\boxed{\mathbf{rate}\mathbf{ }\mathbf{law}\mathbf{=}\mathbf{k}{\left[\mathbf{A}\right]}^{\mathbf{x}}{\left[\mathbf{B}\right]}^{\mathbf{y}}}$

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:

$\mathbf{Rate}\mathbf{ }\mathbf{law}\mathbf{ }\mathbf{=}\mathbf{ }\mathbf{-}\frac{\mathbf{d}\mathbf{\left[}{\mathbf{H}}_{\mathbf{2}}\mathbf{\right]}}{\mathbf{dt}}\mathbf{=}\mathbf{ }{\mathbf{k}}_{\mathbf{2}}\mathbf{\left[}{\mathbf{H}}_{\mathbf{2}}\mathbf{\right]}\mathbf{\left[}{\mathbf{N}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{2}}\mathbf{\right]}\mathbf{ }\mathbf{\left(}\mathbf{1}\mathbf{\right)}$